LIBRARY 

OF  THE 

UNIVERSITY  OF  CALIFORNIA. 


94  J  98 


gale  ^Bicentennial  publication? 

RESEARCH    PAPERS 


FROM  THE 


KENT  CHEMICAL  LABORATORY  OF  YALE 
UNIVERSITY 


gale  ^Bicentennial  publications 

With  the  approval  of  the  President  and  Fellows 
of  TTale  University,  a  series  of  volumes  has  been 
prepared  by  a  number  of  the  Professors  and  In- 
structors, to  be  issued  in  connection  with  the 
Bicentennial  Anniversary,  as  a  partial  indica- 
tion of  the  character  of  the  studies  in  which  the 
University  teachers  are  engaged. 

This   series   of  volumes    is    respectfully  dedicated  to 

of  tt) 


RESEARCH    PAPERS 


FROM  THE 


KENT    CHEMICAL    LABORATORY 


OF 


YALE  UNIVERSITY 


EDITED  BY 

FRANK    AUSTIN    GOOCH 

Professor  of  Chemistry  in  Yale  University 


VOLUME   II. 


NEW   YORK:    CHARLES   SCRIBNER'S   SONS 
LONDON:    EDWARD   ARNOLD 

1901 


Copyright,  1901, 
BY  YALE    UNIVERSITY 

Published,  June^  /go/ 


UNIVERSITY  PRESS    •    JOHN  WILSON 
AND    SON    •    CAMBRIDGE,  U.  S.  A. 


CONTENTS 

VOLUME  II. 

PAGE 
I.    The  Determination  of   Tellurium  by  Precipitation   as  the 

Iodide.    By  F.  A.  GOOCH  and  W.  C.  MORGAN  ....        1 

II.  On  the  Application  of  Certain  Organic  Acids  to  the  Estimation 
of  Vanadium.  By  PHILIP  E.  BROWNING  and  RICHARD 
J.  GOODMAN 4 

III.  The  Determination  of  Oxygen  in  Air  and  in  Aqueous  Solu- 

tion.   By  D.  ALBERT  KREIDER 11 

IV.  A  Method  for  the  Separation  of  Aluminum  from  Iron.    By 

F.  A.  GOOCH  and  F.  S.  HAVENS 20 

V.    The  Estimation   of  Molybdenum  lodometrically.     By  F.  A. 

GOOCH 27 

VI.    The  Application  of  lodic  Acid  to  the  Analysis  of  Iodides.    By 

F.  A.  GOOCH  and  C.  F.  WALKER 33 

VII.    The  Action  of  Urea  and  Primary  Amines  on  Maleic  Anhy- 
dride.   By  FREDERICK  L.  DUNLAP  and  I.  K.  PHELPS    .      42 

VIII.    The  Separation  of  Aluminum  and  Beryllium  by  the  Action  of 

Hydrochloric  Acid.     By  FRANKS  S.  HAVENS    ....      47 

IX.    The  Titration  of  Sodium  Thiosulphate  by  lodic  Acid.     By 

CLAUDE  F.  WALKER 52 

X.    The  Combustion  of  Organic  Substances  in  the  Wet  Way.    By 

I.  K.  PHELPS 62 

XI.    The  Estimation  of  Manganese  as  the  Sulphate  and  as  the 

Oxides.    By  F.  A.  GOOCH  and  MARTHA  AUSTIN  ...       77 
XII.    On  the  Condition  of  Oxidation  of  Manganese  precipitated 
by  the  Chlorate  Process.    By  F.  A.  GOOCH  and  MARTHA 
AUSTIN 85 

XIH.    On  the  Estimation  of  Manganese  separated  as  the  Carbonate. 

By  MARTHA  AUSTIN 96 

XIV.    The  Action  of  Carbon  Dioxide  on  Soluble  Borates.    By  Louis 

CLEVELAND  JONES  .  100 


94198 


viii  CONTENTS 

PAGE 
XV.    Further  Separations  of  Aluminum  by  Hydrochloric  Acid. 

By  FRANKE  STUART  HAVENS 106 

XVI.    The  lodometric  Determination  of  Molybdenum.    By  F.  A. 

GOOCH  and  JOHN  T.  NORTON,  Jr Ill 

XVII.    On  the  Determination  of  Manganese  as   the  Pyrophos- 

phate.    By  F»  A.  GOOCH  and  MARTHA  AUSTIN    .     .     121 

XVIII.  On  the  Detection  of  Sulphides,  Sulphates,  Sulphites,  and 
Thiosulphates  in  the  presence  of  each  other.  By 
PHILIP  E.  BROWNING  and  ERNEST  HOWE  .  .  .  .  134 

XIX.    On  the  Separation  of  Nickel  and  Cobalt  by  Hydrochloric 

Acid.    By  FRANKE  STUART  HAVENS 141 

XX.  The  Ethers  of  Toluquinoneoxime  and  their  bearing  on  the 
Space  Isomerism  of  Nitrogen.  By  JOHN  L.  BRIDGE 
and  WILLIAM  CONGER  MORGAN 145 

XXI.  The  Application  of  Iodine  in  the  Analysis  of  Alkalies  and 
Acids.  By  CLAUDE  F.  WALKER  and  DAVID  H.  M. 
GILLESPIE 162 

XXII.    The  Estimation  of  Boric  Acid.     By  F.  A.  GOOCH  and 

Louis  CLEVELAND  JONES 172 

XXIH.    A  Volumetric  Method  for  the  Estimation  of  Boric  Acid. 

By  Louis  CLEVELAND  JONES 182 

XXIV.    The  Constitution  of  the  Ammonium  Magnesium  Phosphate 

of  Analysis.    By  F.  A.  GOOCH  and  MARTHA  AUSTIN     190 

XXV.  The  Influence  of  Hydrochloric  Acid  in  Titrations  by  So- 
dium Thiosulphate,  with  special  reference  to  the  Esti- 
mation of  Selenious  Acid.  By  JOHN  T.  NORTON,  Jr.  206 
XXVI.  -The  Volatilization  of  the  Iron  Chlorides  in  Analysis,  and 
the  Separation  of  the  Oxides  of  Iron  and  Aluminum. 
By  F.  A.  GOOCH  and  FRANKE  STUART  HAVENS  .  215 

XXVII.  The  Titration  of  Oxalic  Acid  by  Potassium  Permanganate 
in  presence  of  Hydrochloric  Acid.  By  F.  A.  GOOCH 
and  C.  A.  PETERS 222 

XXVIII.  The  Estimation  of  Iron  in  the  Ferric  State  by  Reduction 
with  Sodium  Thiosulphate  and  Titration  with  Iodine. 
By  JOHN  T.  NORTON,  Jr 230 

XXIX.    The  Determination   of   Tellurous  Acid  in  presence  of 

Haloid  Salts.     By  F.  A.  GOOCH  and  C.  A.  PETERS   .     238 

XXX.    An  lodometric  Method  for  the  Estimation  of  Boric  Acid. 

By  Louis  CLEVELAND  JONES 244 


CONTENTS 

XXXI.   The  Double  Ammonium  Phosphates  of  Beryllium,  Zinc, 
and  Cadmium  in  Analysis.     By  MARTHA  AUSTIN 


IX 
PAGE 

252 


XXXII.  Separation  of  Iron  from  Chromium,  Zirconium,  and 
Beryllium,  by  the  Action  of  Gaseous  Hydrochloric 
Acid  on  the  Oxides.  By  FRANKE  STUART  HAVENS 
and  ARTHUR  FITCH  WAY 266 

XXXin.    The  lodometric   Determination  of    Gold.     By  F.  A. 

GOOCH  and  FREDERICK  H.  MORLEY 269 

XXXIV.    The  Action  of  Acetylene  on  the  Oxides  of  Copper.     By 

F.  A.  GOOCH  and  DEFOREST  BALDWIN    .     .     .     .     276 

XXXV.    Notes  on  the  Space  Isomerism  of  the  Toluquinoneoxime 

Ethers.    By  WILLIAM  CONGER  MORGAN  ....    283 

XXXVI.    On  the  Volumetric  Estimation  of  Cerium.    By  PHILIP 

E.  BROWNING 289 

XXXVII.    On  the  Estimation  of  Thallium  as  the  Chromate.     By 

PHILIP  E.  BROWNING  and  GEORGE  P.  HUTCHINS    300 

XXXVIII.  The  Ethics  of  Isonitrosoguaiacol  in  their  relation  to  the 
Space  Isomerism  of  Nitrogen.  By  JOHN  L.  BRIDGE 
and  WILLIAM  CONGER  MORGAN 804 

XXXIX.    The  Constitution  of  the  Ammonium  Magnesium  Arseni- 

ate  of  Analysis.     By  MARTHA  AUSTIN      ....     309 

XL.    On  the  Estimation  of  Thallium  as  the  Acid  and  Neutral 

Sulphates.    By  PHILIP  E.  BROWNING 317 

XLI.    The  Separation  and  Determination  of  Mercury  as  Mer- 

curous  Oxalate.    By  C.  A.  PETERS 320 

XLII.    The  Titration  of  Mercury  by  Sodium  Thiosulphate.    By 

JOHN  T.  NORTON,  Jr 328 

XLIII.    The  lodometric  Estimation  of  Arsenic  Acid.     By  F.  A. 

GOOCH  and  JULIA  C.  MORRIS 336 

XLIV.  On  the  Qualitative  Separation  of  Nickel  from  Cobalt  by 
the  Action  of  Ammonium  Hydroxide  on  the  Ferri- 
cyanides.  By  PHILIP  E.  BROWNING  and  JOHN  B. 
HARTWELL 344 

XLV.  The  Volumetric  Estimation  of  Copper  as  the  Oxalate, 
with  Separation  from  Cadmium,  Arsenic,  Tin,  Iron, 
and  Zinc.  By  CHARLES  A.  PETERS 347 

XL VI.  The  Sulphocyanides  of  Copper  and  Silver  in  Gravi- 
metric Analysis.  By  R.  G.  VAN  NAME  .  .  .  .  359 


x  CONTENTS 

PAGE 

XL VII.  On  the  Estimation  of  Caesium  and  Rubidium  as  the  Acid 
Sulphates,  and  of  Potassium  and  Sodium  as  the  Pyro- 
sulphates.  By  PHILIP  E.  BROWNING 368 

XL VIII.    The  Estimation  of  Calcium,  Strontium,  and  Barium  as  the 

Oxalates.    By  CHARLES  A.  PETERS 373 

XLIX.  The  Action  of  Sodium  Thiosulphate  on  Solutions  of  Me- 
tallic Salts  at  High  Temperatures  and  Pressures.  By 
JOHN  T.  NORTON,  Jr.  .  384 


INDEX  395 


RESEARCH   PAPERS 

PROM  THE 

KENT  CHEMICAL  LABORATORY  OF  YALE 
UNIVERSITY 


OF"  THE 

UNIVERSITY 

OF 


RESEARCH  PAPERS 


THE  DETERMINATION  OF  TELLURIUM  BY 
PRECIPITATION  AS   THE  IODIDE. 

BY  F.  A.  GOOCH  AND  W.  C.  MORGAN.* 

IT  was  known  to  Berzelius  that  hydriodic  acid  and  tellurous 
acid  interact  with  the  formation  of  tellurium  tetraiodide,  which 
is  converted  by  water  into  an  oxyiodide  and  by  excess  of  an 
alkaline  iodide  into  a  soluble  double  salt.  Wheeler  f  has 
shown  that  the  double  salt  which  is  formed  when  tellurous 
iodide  is  boiled  in  a  strong  solution  of  potassium  iodide  in 
dilute  hydriodic  acid  is  definite  and  has  the  constitution 
represented  by  the  formula  2KI  .  TeI4  .  2H2O.  We  have 
observed,  however,  that  when  potassium  iodide  is  added  to  a 
cold  solution  of  tellurous  acid  containing  at  least  one-fourth 
of  its  volume  of  strong  sulphuric  acid,  no  tendency  toward 
the  formation  of  a  double  salt  becomes  apparent  until  the 
potassium  iodide  amounts  to  more  than  enough  to  convert  all 
the  tellurous  acid  present  into  the  tetraiodide  according  to  the 
equation, 

H2Te08  +  4H2S04  +  4KI  =  TeI4  +  4KHS04  +  3H20. 

The  tellurium  tetraiodide  which  is  thus  formed  is  extremely 
insoluble  in  sulphuric  acid  of  the  strength  mentioned,  though 
soluble  in  excess  of  potassium  iodide,  and  acted  upon  by  water 
with  the  formation  of  tellurium  oxyiodide  and  hydriodic  acid. 
It  is  produced  at  first  in  the  condition  of  a  finely  divided 
dark  brown  precipitate  which  upon  agitation  of  the  liquid 
containing  it  gathers  in  curdy  masses  and  settles,  leaving  the 

*  From  Am.  Jour.  Sci.,  ii,  271. 
t  Am.  Jour.  Sci.,  xlv,  267. 
VOL.  n.  —  1 


2  DETERMINATION  OF  TELLURIUM  BY 

supernatant  liquid  clear.  By  taking  advantage  of  this  tendency 
to  curd  it  is  possible  to  determine  without  great  difficulty  the 
exact  point  during  the  gradual  addition  of  potassium  iodide 
when  the  precipitation  of  the  tellurium  iodide  is  complete,  and 
we  have  been  able  to  found  upon  this  property  a  very  simple 
titrimetric  method  for  the  direct  determination  of  small 
amounts  of  tellurium. 

In  our  test  experiments  we  used  tellurium  dioxide  prepared 
by  oxidizing  presumably  pure  tellurium  with  nitric  acid  and 
igniting  the  residue  at  a  low  red  heat.  Weighed  amounts  of 
the  oxide  thus  prepared  were  dissolved  in  Erlenmeyer  beakers 
in  a  very  little  of  a  strong  solution  of  potassium  hydroxide, 
and  dilute  sulphuric  acid  was  added  carefully  until  the  tellu- 
rous  acid  which  was  precipitated  upon  the  neutralization  of 
the  alkaline  hydroxide  was  just  redissolved.  To  this  solution 
sulphuric  acid  of  half-strength  was  added  in  such  amount  that 
the  solution  finally  obtained,  after  adding  the  aqueous  solution 
of  potassium  iodide  subsequently,  should  still  contain  at  least 
one-fourth  of  its  volume  of  strong  sulphuric  acid.  The 
Erlenmeyer  beaker  was  placed  upon  a  pane  of  window  glass 
supported  upon  strips  of  wood  about  1  cm.  above  the  level  of 
the  work  table,  which  was  covered  with  white  paper.  A 
solution  of  approximately  decinormal  potassium  iodide  free 
from  iodate  and  carefully  standardized  in  terms  of  iodine  by  a 
method  described  in  a  former  paper  from  this  laboratory  *  was 
introduced  gradually  from  a  burette  into  the  middle  of  the 
Erlenmeyer  beaker.  As  the  drops  of  the  potassium  iodide 
touched  the  liquid  the  precipitation  formed  at  the  centre  and 
travelled  in  rings  toward  the  outer  walls  of  the  beaker.  When 
the  liquid  became  so  opaque  that  the  effect  of  the  potassium 
iodide  was  distinguished  with  difficulty,  the  beaker  was  rotated 
and  the  curded  precipitate  permitted  to  settle,  and  then  the 
process  of  titration  was  continued  as  before.  We  experimented 
with  amounts  of  tellurium  dioxide  varying  from  approximately 
0.025  grm.  to  0.1  grm.,  the  latter  quantity  being  as  large  as  can 
be  handled  with  accuracy  without  intermediate  removal  of  the 

*  Am.  Jour.  Sci.,  rxxix,  188.     Volume  I,  p.  1. 


PRECIPITATION  AS   THE  IODIDE. 


3 


precipitate  by  filtration.  With  an  Erlenmeyer  beaker  10  cm. 
in  diameter  across  the  bottom  and  a  final  volume  of  liquid 
amounting  to  not  more  than  100  cm3,  we  were  able  to  follow 
the  precipitation  most  easily. 

The  results  of  a  series  of  determinations  made  according  to 
the  method  described  and  recorded  in  the  following  table  are 
closely  accordant,  and,  in  close  agreement  with  the  theory  of 
the  process  if  the  atomic  weight  of  the  tellurium  which  we 
used  is  taken  as  127.  We  feel  justified  in  taking  this  number 
as  the  atomic  weight  of  our  tellurium,  because  the  mean  result 
of  twelve  oxidations  by  standard  potassium  permanganate  of 
tellurium  dioxide,  prepared  similarly  to  that  which  we  used 
and  from  the  same  lot  of  material,  and  the  mean  result  of 
twelve  reductions  by  hydrobromic  acid  of  the  telluric  acid 
thus  produced,*  point  to  this  figure. 


Final 

volume. 

Strongest 
H2S04 

Iodine  value 
of  KI  used. 

TeOj  taken. 

TeO,  found. 

Error. 

present. 

cm8 

cm* 

grm. 

gTm. 

grm. 

grm. 

50 

17 

0.0706 

0.0223 

0.0221 

0.0002- 

50 

17 

0.0764 

0.0244 

0.0239 

0.0005- 

50 

17 

0.1591 

0.0496 

0.0499 

0.0003+ 

60 

17 

0.1655 

0.0517 

0.0519 

0.0002+ 

60 

17 

0.1578 

0.0498 

0.0494 

0.0004- 

80 

30 

0.1591 

0.0498 

0.0499 

0.0001+ 

100 

30 

0.3179 

0.1001 

0.0997 

0.0004- 

100 

30 

0.3186 

0.1008 

0.0999 

0.0009- 

100 

30 

0.3208 

0.1011 

0.1005 

0.0006- 

100 

30 

0.3208 

0.1010 

0.1005 

0.0005- 

From  these  results  it  is  obvious  that  the  method,  which  is 
very  rapid,  is  accurate. 

*  Am.  Jour.  Sci.,  xlviii,  377,  378.    Volume  I,  pp.  279,  281. 


n 


ON   THE   APPLICATION    OF    CERTAIN    ORGANIC 
ACIDS   TO  THE   ESTIMATION  OF  VANADIUM. 

BY  PHILIP  E.  BROWNING  AND  RICHARD  J.  GOODMAN  * 

IN  a  former  paperf  by  one  of  us  a  method  for  the  determi- 
nation of  vanadium  was  described  in  which  tartaric  acid  was 
used  to  reduce  vanadic  acid  to  the  condition  of  the  tetroxide. 
The  method  may  be  briefly  outlined  as  follows : 

Measured  and  weighed  portions  of  a  solution  of  ammonium 
vanadate,  the  standard  of  which  had  been  determined  by  the 
evaporation  and  ignition  of  definite  portions,  were  treated  with 
tartaric  acid  in  excess  and  boiled,  when  the  appearance  of  the 
deep  blue  color  indicated  the  reduction  to  the  condition  of  the 
tetroxide.  After  cooling,  the  solution  was  neutralized  with 
potassium  bicarbonate  and  a  moderate  excess  of  that  reagent 
added.  To  the  alkaline  solution  an  excess  of  a  standard 
solution  of  iodine  was  added  and  the  whole  allowed  to  stand 
about  one  hour,  when  no  further  bleaching  of  the  iodine  was 
noticed.  The  excess  of  iodine  was  then  destroyed  with  a 
standard  solution  of  arsenious  oxide,  starch  was  added,  and  the 
blue  color  obtained  with  a  few  drops  of  the  iodine  solution. 
The  total  amount  of  iodine  used,  less  the  amount  equivalent  to 
the  arsenious  oxide  solution  used,  is  the  amount  necessary  to 
oxidize  the  vanadium  from  the  condition  of  the  tetroxide  to 
that  of  the  pentoxide,  from  which,  according  to  the  following 
equation,  can  be  calculated  the  amount  of  vanadium  present : 
V204  +  I-I  +  H20  =  V206  +  2HI. 

*  From  Am.  Jour.  ScL,  ii,  355. 
t  Zeitschr.  anorg.  Chem.,  vii,  158. 

t  These  determinations  are  best  made  in  small  Erlenmeyer  beakers,  closed 
with  paraffin-coated  corks  while  standing  with  iodine. 


ESTIMATION  OF  VANADIUM. 


The  work  to  be  described  in  this  paper  is  in  part  an 
application  of  the  work  described  in  the  paper  above  mentioned 
to  a  series  of  determinations  of  vanadium  in  the  presence  of 
molybdenum  and  tungsten.  The  solution  of  vanadium  used  was 
one  of  ammonium  vanadate,  and  the  standard  was  determined 
by  evaporating  and  igniting,  in  the  presence  of  a  drop  of 
nitric  acid,  measured  and  weighed  portions,  the  mean  of 
closely  agreeing  results  being  taken  as  the  standard.  Our 
first  series  of  determinations  was  by  the  method  previously 
described,  that  being  the  natural  starting-point  for  the  work 
contemplated.  The  results  follow  in  the  table : 


Bxp. 

V,06  taken. 

V2O8  found. 

Error. 

Tartaric 
acid. 

grm. 

grm. 

grm. 

grm. 

(1) 

0.1621 

0.1618 

0.0003- 

2 

2 

0.1620 

0.1624 

0.0004+ 

2 

(3) 

0.1614 

0.1622 

0.0008+ 

2 

(4) 

0.1619 

0.1606 

0.0013- 

1 

(5) 

0.1604 

0.1597 

0.0007- 

2 

(6) 

0.1618 

0.0615 

00003- 

3 

(7) 

0.1298 

0.1305 

0.0007+ 

1 

(8) 

0.1294 

0.1297 

0.0003+ 

1 

(9) 

0.1618 

0.1618 

O.OOOOi 

2 

(10) 

0.2588 

0.2575 

0.0013— 

3 

(11) 

0.2722 

0.2726 

0.0004+ 

2 

(12) 

0.3273 

0.3269 

0.0004- 

2 

We  next  tried  the  effect  of  treating  a  solution  of  sodium 
tungstate  in  the  same  manner.  We  found  that  after  the 
boiling  with  tartaric  acid,  neutralizing,  adding  iodine  and 
allowing  to  stand  as  before,  the  amount  of  free  iodine  present, 
as  shown  by  the  amount  of  arsenious  oxide  solution  necessary 
to  destroy  it  was  the  same  as  that  originally  added,  showing 
that  no  reduction  had  taken  place.  Accordingly  a  series  of 
determinations  of  vanadium  in  the  presence  of  tungsten  was 
made  which  is  recorded  in  the  next  table. 

The  results  show  that  vanadium  may  be  easily  determined 
by  this  method  in  the  presence  of  tungsten  without  any  evi- 
dent interfering  action  on  the  part  of  the  latter  element. 

When  the  same  method  of  treatment  was  applied  in  the 
presence  of  molybdenum  in  the  form  of  ammonium  molybdate, 


APPLICATION  OF  CERTAIN  ORGANIC  ACIDS 


Exp. 

V2O6  taken. 

V2O5  found. 

Error. 

Sodium 

tuiigstate. 

Tartaric 
acid. 

grm. 

grm. 

grm. 

grm. 

grm. 

(1) 

0.1618 

0.1615 

0.0003- 

1 

3 

(2) 

0.1615 

0.1606 

0.0009- 

1 

3 

(3) 

0.1618 

0.1624 

0.0006+ 

1 

3 

w 

0.1619 

0.1624 

0.0005+ 

1 

3 

w 

0.1627 

0.1623 

0.0004- 

1 

3 

(6) 

0.1621 

0.1624 

0.0003+ 

1 

4 

(7) 

0.2587 

0.2574 

0.0013- 

1 

4 

(8 

0.2587 

0.2689 

0.0002+ 

1 

4 

the  majority  of  the  determinations  gave  large  plus  errors,  and 
a  few  experiments  made  with  the  molybdate  alone  seemed  to 
show  a  noticeable  reduction  of  the  molybdic  acid.  In  the 
following  table  the  results  are  tabulated.  In  experiments  (3), 
(4)  and  (5)  the  mixtures  were  not  boiled  with  tartaric  acid, 
but  warmed  on  a  steam  bath,  with,  however,  no  very  apparent 
prevention  of  the  reducing  action. 


E 

cp. 

V2O6  taken. 

V,O6  found. 

Error. 

Ammonium 
molybdate. 

Tartaric 
acid. 

grm. 

grm. 

grm. 

grm. 

grm. 

1) 

0.1620 

0.1790 

0.0170+ 

2 

2) 

0.1624 

0.1619 

0.0005- 

2 

3 

0.1294 

0.1416 

0.0122+ 

2 

4) 

0.1296 

0.1361 

0.0065+ 

2 

5) 

0.1291 

0.1312 

0.0021+ 

2 

6) 

0.1293 

0.1324 

0.0031+ 

2 

7 

0.1636 

0.1760 

0.0124+ 

2 

8) 

0.1640 

0.1724 

0.0084+ 

2 

[9) 

0.1622 

0.1624 

0.0002+ 

3 

(] 

0) 

0.1622 

0.1632 

0.0010+ 

3 

(1 

1) 

0.1619 

0.1879 

0.0260+ 

3 

(1 

2) 

0.1292 

0.1360 

0.0068+ 

1 

3 

(1 

3) 

0.1860 

0.1917 

0.0057+ 

1 

3 

1 

4) 

0.3274 

0.3733 

0.0459+ 

1 

4 

(1 

6) 

0.2324 

0.2383 

0.0059+ 

1 

4 

We  next  tried  the  action  of  tartaric  acid  upon  the  vanadium 
solution  in  the  cold  and  found  that  the  reduction  could  be 
carried  on  to  completion  under  these  conditions  if  the  tartaric 
acid  was  in  large  excess,  the  time  sufficient,  and  the  volume  of 
the  solution  small.  The  following  series  was  made  to  deter- 
mine these  points : 


TO   THE  ESTIMATION  OF  VANADIUM. 


Exp. 

V2O5  taken. 

V2OS  found. 

Error. 

Time. 

Tartaric 
acid. 

Total 
vol. 

grm. 

grm. 

gnu. 

days. 

grm. 

cms 

(1) 

0.1646 

0.1649 

0.0003+ 

1 

4 

25 

(2) 

0.1640 

0.1606 

0.0034- 

1 

4 

65 

3) 

0.1293 

0.1264 

0.0029— 

2 

3 

55 

4) 

0.1633 

0.1628 

0.0005- 

2 

4 

65 

5) 

0.1293 

0.1288 

0.0005- 

3 

2.5 

60 

6) 

0.1298 

0.1299 

0.0001+ 

3 

2.5 

50 

7) 

0.1295 

0.1279 

0.0016- 

3 

3 

55 

8) 

0.1617 

0.1597 

0.0020- 

4 

2 

70 

9) 

0.1623 

0.1622 

0.0001- 

4 

3 

80 

Solutions  containing  sodium  tungstate  and  ammonium 
molybdate  were  allowed  to  stand  from  one  to  four  days  with 
varying  amounts  of  tartaric  acid  without  giving  any  evidence 
of  reduction. 

In  the  series  which  follows  may  be  seen  the  results  of  a 


Exp. 

V205 
taken. 

V205 
found. 

Error. 

Ammonium 
molybdate. 

Sodium 
tung- 

state. 

Time 
in 
days. 

Total 
volume. 

Tartaric 
acid. 

grm. 

grin. 

grin. 

grm. 

grm. 

cms 

grm. 

(1) 

0.1552 

0.1558 

0.0006+ 

1 

25 

5 

(2) 

0.1289 

0.1301 

0.0012-f 

< 

1 

25 

5 

(3) 

0.2583 

0.2587 

0.0004+ 

m 

§ 

1 

50 

5 

(4) 

0.1293 

0.1299 

0.0006+ 

1 

1 

25 

6 

(5) 

0.2582 

0.2591 

0.0009+ 

1 

t 

1 

50 

6 

(6) 

0.2582 

0.2588 

0.0006+ 

1 

1 

50 

5 

(7) 

0.1297 

0.1308 

0.0011+ 

1 

1 

25 

5 

(8) 

0.1291 

0.1289 

0.0002- 

1 

1 

25 

6 

(9) 

0.2582 

0.2568 

0.0014- 

e 

1 

1 

50 

5 

(10) 

0.1293 

0.1299 

0.0006+ 

i' 

1 

1 

25 

8 

(11) 

0.2582 

0.2579 

0.0003- 

i 

1 

1 

50 

5 

(12) 

0.1550 

0.1538 

0.0012- 

. 

t 

2 

25 

5 

(13) 

0.1556 

0.1545 

0.0011- 

e 

f 

2 

25 

5 

(14) 

0.1289 

0.1296 

0.0007+ 

t 

f 

2 

25 

5 

(15) 

0.1549 

0.1527 

0.0022- 

0.5 

2 

25 

5 

(16) 

0.1553 

0.1548 

0.0005- 

1 

m 

2 

25 

5 

(17) 

0.1556 

0.1554 

0.0002- 

1 

2 

25 

5 

(18) 

0.1293 

0.1310 

0.0017+ 

1 

> 

2 

25 

6 

(19) 

0.1295 

0.1299 

0.0004+ 

. 

y 

2 

25 

6 

(20) 

0.1293 

0.1289 

0.0004- 

r 

i 

2 

25 

7 

(21) 

0.1293 

0.1301 

0.0008+ 

( 

3 

25 

5 

(22) 

0.1289 

0.1299 

0.0010+ 

0.5 

t 

3 

25 

5 

(23) 

0.1293 

0.1292 

0.0001- 

1 

f 

3 

25 

7 

(24) 

0.1556 

0.1567 

0.0011+ 

1 

t 

3 

30 

5 

(25) 

0.1291 

0.1289 

0.0002- 

1 

i 

3 

25 

7 

(26) 

0.1550 

0.1557 

0.0007+ 

. 

^ 

4 

25 

5 

(27) 

0.1554 

0.1557 

0.0003+ 

1 

t 

4 

25 

5 

(28) 

0.1556 

0.1557 

0.0001+ 

0.5 

•  • 

4 

25 

6 

8  APPLICATION  OF  CERTAIN  ORGANIC  ACIDS 

number  of  determinations  of  vanadium  in  the  presence  of 
molybdenum  and  tungsten  made  in  the  cold  and  allowed  to 
stand  from  one*  to  four  days.  It  will  be  noticed  that  the 
results  on  standing  one  day  with  five  grams  of  tartaric  acid 
are  for  the  most  part  satisfactory,  and  an  increase  in  the 
length  of  tune  does  not  cause  any  apparent  reduction  of 
the  molybdenum. 

Friedheim  f  has  shown  that  vanadium  is  reduced  from  the 
condition  of  the  pentoxide  to  that  of  the  tetroxide  by  boiling 
with  oxalic  acid.  The  reduction  is  so  complete  that  he  has 
developed  a  method  for  the  estimation  of  vanadium  upon  this 
reaction  and  shows  that  it  may  be  applied  in  the  presence  of 
molybdenum  and  tungsten,  the  acids  of  these  elements  not 
being  reduced  by  the  oxalic  acid.  When  the  vanadic  acid  is 
reduced  the  oxalic  acid  is  oxidized  and  a  definite  amount  of 
carbon  dioxide  evolved  according  to  the  reaction. 

V206  +  H2C204  =  V204  +  H20  +  2C02. 

This  carbon  dioxide  Friedheim  conducts  by  an  appropriate 
form  of  apparatus  into  potassium  hydroxide  and  weighs. 
From  this  weight  the  amount  of  vanadic  acid  originally 
present  may  be  readily  calculated. 

We  have  applied  the  method  of  oxidation  with  standard 
iodine  described  in  the  tartaric  acid  process  to  the  residue 
after  boiling  with  oxalic  acid.  The  method  of  treatment  was 
identical  with  that  outlined  at  the  beginning  of  this  paper. 
The  results  which  follow  in  the  table  are  for  the  most  part 
satisfactory  and  the  method  is  certainly  more  easily  applied 
than  Friedheim's  process,  the  potassium  hydroxide  absorption 
apparatus  being  unnecessary. 

Having  applied  successfully  both  tartaric  acid  and  oxalic 
acid  in  the  manner  described,  the  action  of  citric  acid  applied 
in  the  same  manner  suggested  itself  as  a  fitting  conclusion  to 
the  study  of  the  action  of  this  class  of  organic  acids.  In  this 

*  Some  of  the  determinations  designated  in  the  table  as  having  stood  one 
day  in  reality  stood  only  about  fifteen  hours,  from  6  p.  M.  to  9  A.  M. 
t  Zeitschr.  anorg.  Chem.,  i,  312. 


TO   THE  ESTIMATION  OF  VANADIUM. 


Eip. 

V,O6  taken. 

V,O6  found. 

Error. 

Oxalic 
acid. 

Ammonium 
molybdate. 

Sodium 
tungstate. 

gnu. 

gnu. 

gnu. 

gnu. 

gnu. 

gnu. 

(1) 

0.1806 

0.1803 

0.0003- 

1 

2 

0.1950 

0.1955 

0.0005+ 

1 

3) 

0.1959 

0.1955 

0.0004- 

1 

, 

(4) 

0.1950 

0.1959 

0.0009+ 

1 

t 

(5) 

0.1954 

0.1977 

0.0023+ 

1 

t 

(6) 

0.1956 

0.1960 

0.0004+ 

1 

m 

(7) 

0.1956 

0.1964 

0.0008+ 

1 

(8) 

0.1956 

0.1957 

0.0001+ 

1 

(9) 

0.3900 

0.3899 

0.0001- 

2 

(10) 

0.3897 

0.3917 

0.0020+ 

2 

t 

(11) 

0.3903 

0.3905 

0.0002+ 

2 

(12) 

0.1954 

0.1959 

0.0005+ 

2 

Y 

(13) 

0.1957 

0.1960 

0.0003+ 

2 

i 

(14) 

0.1954 

0.1961 

0.0007+ 

2 

i 

(15) 

0.1806 

0.1818 

0.0012+ 

3 

e 

(16) 

0.1807 

0.1827 

0.0020+ 

3 

(17) 

0.1809 

0.1803 

0.0006- 

3 

Y 

(18) 

0.1956 

0.1961 

0.0005+ 

3 

t 

1 

(19) 

0.3611 

0.3617 

0.0006+ 

5 

, 

(20) 

0.3616 

0.3626 

0.0010+ 

5 

i 

case,  as  in  the  others,  the  reduction  of  the  vanadic  acid  is 
easily  and  quickly  effected,  but  the  oxidation  with  the  iodine 
is  slower  than  in  the  presence  of  alkaline  oxalates  and  tar- 
trates.  In  the  case  of  oxalic  acid  on  standing  about  fifteen 
minutes  with  the  excess  of  iodine,  when  tartaric  acid  has 
been  used,  the  bleaching  of  the  iodine  continues  from  thirty  to 
forty-five  minutes,  but  in  the  presence  of  the  alkaline  citrate 
the  time  required  is  fully  an  hour.  A  large  excess  of  tartaric 
or  oxalic  acids  does  not  seem  to  materially  affect  the  results, 
but  hi  the  use  of  citric  acid  it  is  advisable  to  avoid  a  large 
excess,  which  tends  to  give  high  results.  Accordingly  in  the 
following  series  of  experiments  it  will  be  noticed  that  the 
amounts  of  citric  acid  do  not  exceed  two  grams  except  where 
ammonium  molybdate  or  sodium  tungstate  are  present,  when 
the  ammonium  or  sodium  base  combines  with  part  of  the  free 
acid.  The  results  follow,  on  p.  10. 

The  mode  of  proceeding  in  the  estimation  of  vanadium  by 
the  use  of  either  tartaric,  oxalic,  or  citric  acid  may  be  briefly 
summarized  as  follows :  To  a  solution  of  a  vanadate  with  or 
without  a  tungstate  or  molybdate,  add  approximately  one 


10 


ESTIMATION  OF  VANADIUM. 


E 

*I 

V,O4  taken. 

V2O6  found. 

Error. 

Citric 
acid. 

Ammonium 
molybdate. 

Sodium 
tungstate. 

gnn. 

grin. 

gmi. 

gnn. 

grm. 

gnn. 

(1 

i 

0.1956 

0.1956 

0.0000 

1 

2 

0.3905 

0.3921 

0.0016+ 

2 

3 

0.1960 

0.1960 

0.0000 

1 

4 

0.1953 

0.1960 

0.0007+ 

1 

5 

0.2088 

0.2082 

0.0006- 

2 

6 

0.2100 

0.2098 

0.0002- 

2 

0.2092 

0.2107 

0.0015+ 

1 

8 

i 

0.2092 

0.2107 

0.0015+ 

2 

9 

i 

0.2096 

0.2082 

0.0014- 

2 

0.5 

(1 

0 

i 

0.2099 

0.2116 

0.0017+ 

3 

0.5 

1 

0.2095 

0.2101 

0.0006+ 

2 

6.5 

(1 

2 

0.2099 

0.2095 

0.0004- 

3 

.0.6 

gram  of  the  acid  for  every  tenth  of  a  gram  of  substance  to  be 
determined.  Heat  the  solution  to  boiling,  except  in  case  tar- 
taric  acid  be  present  with  molybdic  acid,  when  digestion  for 
from  fifteen  to  twenty-four  hours  in  the  cold  should  be  substi- 
tuted. To  the  cold  liquid  add  about  five  grams  of  potassium 
bicarbonate  for  every  gram  of  acid  used.  Add  iodine  in  slight 
excess  and  set  aside  until  no  further  bleaching  is  noticeable. 
Destroy  the  excess  of  iodine  with  arsenious  oxide  solution,  add 
starch,  and  titrate  back  with  standard  iodine.  The  total 
amount  of  iodine  used  less  the  equivalent  of  the  arsenious 
oxide  is  the  measure  of  the  oxidation. 

We  have  found  it  of  advantage,  when  starting  with  a  new 
solution  of  the  vanadate,  to  make  one  determination  roughly 
and  to  get  from  this  rough  determination  the  proportions  of 
acid  and  iodine  to  be  used  in  the  determinations  to  follow. 
Large  amounts  of  the  acid  and  a  large  excess  of  the  iodine 
have  been  employed  in  many  determinations  without  any 
apparent  unfavorable  effect  upon  the  results.  The  tendency, 
however,  under  these  circumstances  is  toward  plus  errors, 
which  may  be  avoided  by  following  the  above  directions. 


Ill 


THE  DETERMINATION  OF  OXYGEN  IN  AIR  AND 
IN  AQUEOUS   SOLUTION. 

BY  D.  ALBERT  KREIDER.* 

WHILE  there  is  little  to  be  hoped  for  by  way  of  improve- 
ment in  the  accuracy  of  present  known  methods  for  the  deter- 
mination of  oxygen  in  the  air,  some  choice  as  to  manipulation 
may  nevertheless  be  desirable,  and  a  process  which  is  not  lim- 
ited wholly  to  the  methods  and  apparatus  of  ordinary  gas 
analysis  will  doubtless  often  be  found  serviceable. 

The  very  satisfactory  results  which  I  have  obtained  in  the 
determination  of  perchlorates  by  the  action  of  the  liberated 
oxygen  upon  hydriodic  acid  through  the  medium  of  nitric 
oxidef  has  led  me  to  test  this  action  upon  the  oxygen  of  the 
air,  where  only  the  smaller  amount  of  oxygen  and  its  greater 
dilution  with  nitrogen  might  be  expected  to  be  unfavorable. 
However,  with  the  apparatus  and  manipulation  herein  de- 
scribed it  will  be  seen  that  the  method  affords  a  means  for  the 
determination  of  the  oxygen  of  the  air  or  of  dissolved  oxygen 
with  ease  and  rapidity  and  with  sufficient  accuracy  for  all 
practical  purposes. 

The  method  in  brief  is  simply  the  conducting  of  a  known 
volume  of  air  through  a  strong  solution  of  hydriodic  acid  in 
the  presence  of  nitric  oxide ;  subsequently  neutralizing  the 
acid  with  potassium  bicarbonate  and  titrating  the  liberated 
iodine  with  standard  decinormal  arsenic  solution  from  which 
the  equivalent  volume  of  oxygen  is  readily  calculated.  By 
several  simple  devices,  to  be  described,  all  calculations  may  be 
done  away  with  and  the  percentage  of  oxygen  seen  imme- 

*  From  Am.  Jour.  Sci.,  ii,  361. 

t  Am.  Jour.  Sci.,  1,  287.    Volume  I,  p.  316. 


12  THE  DETERMINATION  OF  OXYGEN 

diately  by  the  volume  of  arsenic  solution  required  for  the 
titration. 

The  volume  of  oxygen  found  by  means  of  the  arsenic 
solution  is,  of  course,  under  the  standard  conditions  of  tem- 
perature and  pressure  (0°  and  760  mm.),  and  it  is  therefore 
essential  either  to  calculate  this  volume  into  that  which  it  would 
occupy  under  the  conditions  of  the  experiment  or  to  reduce 
to  the  standard  conditions  of  temperature  and  pressure  the 
volume  of  air  taken.  The  latter  plan  is  the  more  satisfactory 
since  by  Lunge's  ingenious  device*  the  reduction  can  be 
readily  effected  without  any  calculation  and  independently  of 
changing  temperature  and  pressure.  For  my  purpose  the 
following  arrangement  of  two  burettes  answered  admirably. 
One  burette  graduated  to  120  cm3  contained  over  mercury 
the  same  volume  of  moist  air  which  100  cm3  of  air  at  0°  and 
760  mm.  would  occupy  under  the  given  conditions,  this  stand- 
ard being  very  carefully  determined.  By  means  of  a  T-tube 
this  standard  burette  was  placed  between  and  in  connection 
with  the  burette  in  which  the  volume  of  air  to  be  analyzed 
was  measured,  and  a  movable  reservoir  of  mercury.  Both 
burettes  were  firmly  fastened  to  a  movable  iron  rod  and  the 
zero  marks  accurately  adjusted  to  the  same  level.  By  draw- 
ing into  the  measuring  burette  a  volume  of  air  greater  than 
that  required  —  for  which  purpose  a  small  bulb  was  attached 
to  the  lower  end  of  the  burette,  and  then  by  raising  the  reser- 
voir of  mercury,  compressing  the  air  in  the  standard  tube  to 
the  100  cm3  mark,  at  the  same  time  allowing  the  excess  of 
air  to  escape  from  the  measuring  burette,  exactly  100  cm3  of 
air  under  the  standard  conditions  of  temperature  and  pressure 
was  obtained.  To  facilitate  the  adjustment,  two  strips 
of  wood  were  fastened  to  the  rubber  connection  by  means  of 
screw  pinch-cocks  in  such  a  way  that  by  closing  one  pinch- 
cock  the  flow  of  mercury  from  the  reservoir  could  be  shut  off, 
and  then  by  gradually  tightening  the  other  pinch-cock  the 
mercury  would  be  forced  out  of  the  rubber,  and  thus  an  easy 
and  accurate  adjustment  to  the  100  cm3  mark  be  secured. 
*  Zeitschr.  angew.  Chem.,  1890, 139. 


IN  AIR  AND  IN  AQUEOUS  SOLUTION.  13 

The  apparatus  in  which  the  action  of  the  oxygen  upon 
hydriodic  acid  was  effected  consisted  of  a  300  cm3  bulb  pipette, 
both  ends  of  which  were  cut  off  short  and  sealed  to  glass 
stop-cocks.  The  tube  from  one  of  the  stop-cocks  was  cut  off 
short  after  being  tapered  and  constricted  so  as  to  hold  a  rubber 
connector  tightly,  while  the  tube  from  the  other  stop-cock 
was  left  sufficiently  long  to  reach  to  the  bottom  of  a  500  cm8 
Erlenmeyer  beaker.  These  tubes  are  preferably  of  about 
3  mm.  bore,  since  for  the  several  connections  all  air  may  be 
expelled  from  tubes  of  this  size  by  displacement  with  water. 
In  order  to  expel  all  air  from  the  flask,  instead  of  passing  a 
current  of  carbon  dioxide  as  was  done  in  the  determination 
of  perchlorates,  tune  was  saved  by  first  filling  the  flask  with 
water,  which  was  then  displaced  by  pure  carbon  dioxide 
(prepared  as  described  below)  and  the  flask  subsequently 
exhausted,  which  was  accomplished  instantaneously  by  the 
device  described  in  the  article  on  perchlorates.  The  required 
amounts  of  potassium  iodide  solution,  hydrochloric  acid  and 
nitric  oxide  were  drawn  in  in  the  order  named,  after  which 
the  measured  volume  of  air  was  gradually  admitted  Awhile 
the  bulb  was  constantly  agitated  so  as  to  keep  the  hydriodic 
acid  continually  renewed  along  the  surface  of  the  bulb. 
The  shaking  was  continued  for  a  minute  or  two  until  the 
action  was  completed,  when  a  dilute  solution  of  potassium 
bicarbonate  was  admitted.  The  carbon  dioxide  liberated 
forces  the  liquid  from  the  bulb  into  a  beaker  which  contains 
bicarbonate  in  amount  sufficient,  as  previously  determined,  to 
neutralize  all  the  acid  taken.  When  the  exit  is  too  slow  more 
bicarbonate  may  be  admitted  through  the  other  stop-cock,  and 
after  neutralization  has  been  completed  the  bulb  may  be 
washed  out  without  any  danger  from  the  admission  of  air. 

All  the  water  employed,  both  for  the  solution  of  potassium 
iodide  and  for  the  various  connections,  was  free  of  oxygen. 
It  was  prepared  by  filling  a  three-liter  flask  with  distilled 
water  and  boiling  until  the  volume  of  the  liquid  was  reduced 
about  one-third,  when  the  flask  was  closed  by  a  doubly 
perforated  rubber  stopper  and  fitted  as  a  wash  bottle.  By 


14  THE  DETERMINATION  OF  OXYGEN 

means  of  the  tube  which  reached  below  the  surface  of  the 
water,  pure  carbon  dioxide  was  passed  through  while  the 
water  was  still  boiling,  which  together  with  the  escaping 
steam  was  sure  to  expel  all  oxygen.  Then  the  heat  was 
removed  and  the  current  of  carbon  dioxide  continued  until 
the  boiling  ceased,  when  the  escape  tube  was  closed  by  a 
piece  of  rubber  tubing  and  screw  pinch-cock.  As  the  water 
cooled  it  was  well  shaken  while  still  in  connection  with 
the  carbon  dioxide  generator,  and  thus  became  saturated 
with  the  gas,  which  was  then  pumped  in  under  considerable 
pressure  by  the  little  hand  pump  described  in  a  previous 
paper  from  this  laboratory.  By  this  means  the  water  could 
be  drawn  as  needed  without  the  introduction  of  any  air. 
The  escape  tube  was  provided  with  a  rubber  tube  and  screw 
pinch-cock,  and  a  long,  slender  nozzle  which  could  be  in- 
serted into  the  tubes  of  the  absorption  apparatus.  A  bottle 
thus  charged  sufficed  for  all  the  determinations  and  required 
only  an  occasional  supply  of  carbon  dioxide  when  large 
draughts  of  water  were  required  for  making  the  potassium 
iodide  solution. 

The  potassium  iodide  solution  was  made  up  to  contain  one 
gram  of  the  salt  in  thirty  cubic  centimeters  of  water,  and 
was  contained  in  convenient  form  in  an  ordinary  wide- 
mouthed  bottle  fitted  as  a  wash  bottle,  and  graduated  ap- 
proximately for  each  thirty  cubic  centimeters'  volume  —  the 
amount  usually  taken.  The  potassium  iodide  was  weighed 
into  the  bottle,  which  was  then  closed  and  all  air  expelled 
by  a  current  of  carbon  dioxide,  when  the  desired  amount  of 
water,  free  of  oxygen,  was  drawn  in,  and  attachment  again 
made  with  the  carbon  dioxide  generator.  After  allowing  the 
gas  to  pass  for  several  minutes  the  exit  was  closed  and  the 
gas  pumped  in  by  the  little  hand  pump.  Inasmuch  as  this 
solution,  when  it  was  used,  was  drawn  into  an  exhausted 
bulb,  the  bottle  could  be  emptied  without  ever  exposing  its 
contents  to  the  air. 

Nitric  oxide  was  generated  very  satisfactorily  according  to 
Professor  Gooch's  method  —  by  the  action  of  nitric  acid 


IN  AIR  AND  IN  AQUEOUS  SOLUTION.  15 

upon  globules  of  copper  in  a  Kipp  generator.  When  the 
nitric  acid  is  diluted  with  an  equal  volume  of  water  the 
evolution  of  the  gas  is  sufficiently  rapid  without  the  applica- 
tion of  heat,  but  contamination  by  the  higher  oxide  is  more 
likely.  However,  since  it  is  necessary,  in  order  to  be  certain 
of  purity,  to  pass  the  gas  through  an  acidified  solution  of 
potassium  iodide  before  applying  it  to  the  determination 
of  oxygen,  whatever  higher  oxide  may  be  present  will  be 
reduced.  By  passing  the  gas,  as  it  issued  from  the  generator, 
through  a  set  of  Geisler  bulbs  containing  an  acidified  solution 
of  potassium  iodide  and  washing  with  potassium  iodide  solu- 
tion, the  perfectly  purified  gas  was  obtained.  Theoretically, 
only  a  small  amount  of  the  nitric  oxide  is  required  for  the 
transference  of  the  oxygen  to  the  hydriodic  acid,  but  when 
too  little  is  taken  the  action  is  very  slow.  On  the  other 
hand,  too  large  an  amount  relieves  the  vacuum  to  such  an 
extent  as  to  interfere  with  the  introduction  of  the  air.  A 
little  device  to  measure  the  volume  of  gas  taken  was  there- 
fore attached  to  the  generator.  It  consisted  of  a  tube  filled 
with  water  and  roughly  graduated  for  every  five  cubic  centi- 
meters, so  attached  to  the  generator  that  the  gas  would  enter 
by  displacement  of  the  water,  which  would  descend  to  a 
lower  bulb,  and  as  the  gas  was  withdrawn  the  water  would 
again  take  its  place.  Fifteen  cubic  centimeters  of  the  gas 
was  found  a  convenient  and  satisfactory  amount  for  the 
analysis. 

Carbon  dioxide  was  generated  in  a  Kipp  generator,  the 
acid  and  marble  of  which  had  been  previously  boiled  and 
contained  a  little  cuprous  chloride.  To  remove  a  trace  of 
reducing  matter  which  the  gas  was  found  to  contain,  it  was 
first  passed  through  a  solution  of  iodine  and  washed  with 
potassium  iodide. 

For  the  titration  a  decinormal  solution  of  arsenious  oxide 
(4.95  grms.  to  the  liter)  was  employed :  1  cubic  centimeter 
being  equal  to  0.559846  cm3  of  oxygen  at  0°  and  760  mm. 
when  the  weight  of  a  liter  of  oxygen  at  0°  and  760  mm.  is 
taken  as  1.42895  grin.  When  the  volume  of  air  taken 


16 


THE  DETERMINATION  OF  OXYGEN 


is  100  cm3  under  standard  conditions  of  temperature  and 
pressure,  as  obtained  by  Lunge's  device,  the  following  table, 
calculated  for  the  volume  of  oxygen  equivalent  to  the  volume 
of  arsenic  solution,  shows  directly  the  percentage  of  oxygen 
corresponding  to  the  reading  of  the  burette.  The  correction 
necessary  for  the  fraction  of  a  tenth  of  a  cubic  centimeter 
of  the  arsenic  solution  is  obtained  with  sufficient  accuracy 
by  simply  multiplying  by  0.005. 

RELATION  OF  ARSENIC  TO  OXYGEN. 


Correction  for 

SKA. 

Oxygen  equivalent 
at  0°  and  760  mm. 

0.01  cms—  AsjjO,, 

cms 

37.0 

cm3 
20.714 

0.005 

37.1 

20.770 

37.2 

20.826 

37.3 

20.882 

37.4 

20.938 

37.5 

20.994 

37.6 

21.050 

37.7 

21.106 

37.8 

21.162 

37.9 

21.218 

38.0 

21.274 

Table  I  shows  the  results  obtained  in  a  series  of  determina- 
tions. Experiments  (1)  to  (11)  inclusive  were  made  upon  por- 
tions of  air  collected  over  water  on  March  28,  measured  in  an 
ordinary  gas  burette  and  reduced  to  the  standard  conditions 
of  temperature  and  pressure.  The  remainder  of  the  deter- 
minations were  made  upon  air  collected  on  April  8,  each 
portion  having  been  measured  in  the  apparatus  described,  for 
the  reduction  to  standard  conditions. 

No  correction  was  found  necessary  for  the  blank  determina- 
tions, since  when  boiled  water  was  used  the  solution  was  only 
faintly  colored  with  iodine,  which  requires  only  a  drop  of 
arsenic  solution  to  bleach  it.  As  is  evident  from  the  table,  the 
determinations  according  to  this  method  are  not  reliable  beyond 
0.05  per  cent,  but  for  practical  purposes  this  is  sufficiently 
accurate.  For  the  sake  of  comparison  two  determinations  by 


IN  AIR  AND  IN  AQUEOUS  SOLUTION. 
TABLE  I. 


17 


Kxp. 

Volume  of  Air 
reduced  to  0° 
and  760  mm. 

l>°" 
required. 

Volume  of 
Oxygen  found  at 
0°  and  760  mm. 

Per  cent  of 
Oxygen  in  Air. 

cm3 

cms 

cnjS 

(1) 

91.18 

34.06 

19.07 

20.91 

(2) 

91.73 

34.47 

19.30 

21.04 

(3) 

90.84 

34.25 

19.17 

21.11 

(4) 

90.60 

34.20 

19.16 

21.13 

(5) 

86.06 

32.55 

18.22 

21.17 

(6) 

85.96 

32.40 

18.14 

21.10 

(7) 

86.49 

32.53 

18.21 

21.06 

(8) 

87.85 

33.00 

18.47 

21.03 

(9) 

44.17 

16.60 

9.29 

21.04 

(10) 

44.11 

16.70 

9.35 

21.19 

(11) 

44.54 

16.80 

9.41 

21.12 

(12) 

100.00 

37.44 

20.96 

20.96 

(13) 

100.00 

37.54 

21.01 

21.01 

(14) 

100.00 

37.50 

20.99 

20.99 

(15) 

100.00 

37.57 

21.03 

21.03 

(16) 

100.00 

37.47 

20.97 

20.97 

(17) 

100.00 

37.50 

20.99 

20.99 

the  pyrogallic  acid  method  were  made  upon  a  portion  of  the 
same  air  used  in  the  last  experiments,  the  results  being  20.93 
and  20.88  per  cent  respectively.  While  the  pyrogallic  acid 
method  is  capable  of  much  greater  accuracy  when  applied  in 
Hempel's  improved  apparatus,  in  ordinary  burettes  it  will 
probably  not  yield  more  closely  agreeing  results  than  the 
above  method. 

Determination  of  dissolved  Oxygen.  —  A  deter- 
mination of  oxygen  dissolved  in  water  can  be 
completed  by  the  above  method  in  about  ten 
minutes  by  means  of  the  apparatus  illustrated  by 
the  accompanying  figure. 

The  apparatus  consisted  of  a  flask  of  about  300 
cm3  capacity,  into  the  bottom  of  which  was  sealed 
a  stop-cock  with  a  long  exit  tube.  Upon  the 
neck  was  cut  the  fiducial  circle  c  and  immediately 
above  this  stop-cock  e  was  sealed  as  shown.  The 
neck  of  the  flask  was  drawn  out  and  sealed  to 
stop-cock  d  and  the  bulb,  a,  of  about  30  cm3 
capacity  blown  in  it.  The  capacity  of  the  apparatus  be- 

VOL.   II. 2 


FIG.  17. 


18  THE  DETERMINATION  OF  OXYGEN 

tween  stop-cock,  &,  and  the  fiducial  mark,  c,  was  carefully 
determined. 

The  manipulation  for  the  determination  of  dissolved  oxygen 
was  as  follows :  The  flask  was  held  in  the  position  shown  by 
a  clamp  fastened  to  a  movable  support.  Stop-cock  b  being 
closed,  the  water  was  admitted  through  e  and  the  air  allowed 
to  escape  through  d  until  the  level  of  water  was  that  indicated 
by  the  line  /.  (When  the  water  to  be  examined  is  not 
saturated  with  air,  the  flask  must  first  be  filled  with  carbon 
dioxide  and  the  water  entered  by  replacement  of  that  gas.) 
With  d  closed,  sufficient  water  was  allowed  to  escape  through 
b  to  bring  the  surface  to  e,  which  was  then  closed.  The  nitric 
oxide  generator  was  then  attached  to  d,  and  by  opening  b  the 
gas  was  allowed  to  replace  the  water  until  the  meniscus 
coincided  with  c,  when  d  was  closed  and  the  generator 
disconnected.  Two  cubic  centimeters  of  strong  hydrochloric 
were  introduced  through  e  by  expelling  nitric  oxide  through 
d,  in  which  a  drop  of  water  formed  an  effective  trap  to  prevent 
the  entrance  of  air.  Then  the  potassium  iodide  was  admitted 
in  the  same  way.  The  solution  of  iodide  for  this  purpose 
was  free  of  oxygen  and  contained  one  gram  in  three  cubic 
centimeters.  It  was  kept  under  pressure  of  carbon  dioxide  in 
the  bottle  previously  described,  and  by  means  of  a  long  nozzle 
could  be  conducted  to  the  bottom  of  eh  and  thus  be  admitted 
with  but  momentary  and  slight  contact  with  the  air.  The 
tube  eh  contained  approximately  three  cubic  centimeters. 
With  all  the  stop-cocks  closed,  the  flask  was  inverted  several 
times  and  thoroughly  shaken,  at  the  same  tune  washing  out 
the  ends  of  the  stop-cocks  with  distilled  water.  After  again 
placing  the  apparatus  in  its  position,  enough  potassium  bicar- 
bonate solution  was  admitted  through  e  to  expel  all  the 
nitric  oxide  through  d\  the  bulb,  a,  holding  sufficient  of  the 
bicarbonate  to  neutralize  all  the  acid  taken.  The  bicarbonate 
being  heavier  quickly  diffuses  through  the  contents  of  the 
flask  and  neutralizes  the  acid;  d  and  e  are  kept  closed  for  a 
minute  with  b  open  so  as  to  allow  sufficient  of  the  liquid  to 
escape  into  a  beaker  containing  some  bicarbonate  to  provide 


IN  AIR  AND  IN  AQUEOUS  SOLUTION. 


19 


space  for  the  carbon  dioxide  evolved.  Then  the  flask  is  washed 
out  and  its  contents  titrated  with  arsenic. 

The  bleaching,  by  the  aid  of  starch  for  the  final  reaction, 
can  be  accurately  read  to  a  single  drop  Usually  the  reading 
was  verified  by  adding  a  drop  of  £$  iodine  solution,  which 
produced  the  characteristic  color. 

Table  II  gives  the  results  of  a  series  of  determinations. 

TABLE  IL 


Volume  of  Water 
taken. 

Temperature. 

As203  required. 

Volume  of  Oxygen 
dissolved  in  lOOOcmS 
of  water  at  760  mm. 

cm* 

°C. 

cm8 

cm3 

314.63 

20 

3.42 

6.04 

314.63 

20 

3.45 

6.09 

314.63 

20 

3.40 

6.00 

314.63 

20 

3.41 

6.02 

314.63 

20 

3.43 

6.05 

314.63 

20 

3.40 

6.00 

314.63 

20 

3.36 

5.93 

314.63 

20 

3.40 

6.00 

314.63 

20 

3.40 

6.00 

314.63 

20 

3.50 

6.18 

314.63 

20 

3.38 

5.96 

314.63 

20 

3.40 

6.00 

The  mean  of  these  determinations  gives  6.022  cm3  of  oxygen 
as  the  amount  dissolved  in  distilled  water  at  20°  C.  and  760 
mm.,  and  while  some  of  the  determinations  vary  considerably 
from  this  mean,  as  a  whole  they  are  fairly  accordant.  This 
method,  moreover,  is  applicable  to  carbonated  water. 


IV 


A  METHOD  FOR  THE  SEPARATION  OF 
ALUMINUM  FROM  IRON. 

BY  F.  A.  GOOCH  AND  F.  S.  HAVENS.* 

OF  the  well-known  methods  for  the  separation  of  aluminum 
from  iron  —  by  the  action,  for  example,  of  an  alkaline  hydroxide 
in  aqueous  solution  or  by  fusion  of  the  mixed  oxides  in 
potassium  or  sodium  hydroxide ;  by  reduction  of  the  iron 
oxide  to  the  metal  by  heating  in  hydrogen,  with  the 
subsequent  solution  of  the  metallic  iron  in  hydrochloric 
acid;  by  boiling  the  nearly  neutral  solution  of  the  salts  of 
aluminum  and  iron  with  sodium  thiosulphate  either  with  or 
without  sodium  phosphate  ;  by  acting  with  hydrogen  sulphide 
or  ammonium  sulphide  upon  solutions  of  the  salts  containing 
also  an  ammoniacal  citrate  or  tartrate  —  no  single  process  can 
be  said  to  be  ideal  as  regards  directness,  rapidity  and  accuracy 
of  working.  We  have  deemed  it  not  superfluous,  therefore, 
to  attempt  the  utilization  of  a  reaction  which  should  apparently 
be  capable  of  effecting  directly  and  quickly  the  separation  of 
aluminum  from  iron  under  conditions  easily  attainable. 

It  is  known  t  that  the  hydrous  aluminum  chloride  A1C13.6H2O 
is  very  slightly  soluble  in  strong  hydrochloric  acid,  while  ferric 
chloride,  on  the  other  hand,  is  extremely  soluble  in  that 
medium.  It  is  this  difference  of  relation  of  which  we  wished 
to  take  advantage. 

It  appeared  at  the  outset  that  crude  aluminum  chloride 
could  be  freed  from  every  trace  of  a  ferric  salt  by  dissolving  it 
in  the  least  possible  amount  of  water,  saturating  the  cooled 
solution  with  gaseous  hydrochloric  acid,  filtering  upon  asbestos 

*  From  Am.  Jour.  Sci.,  ii,  416. 

t  Gladysz,  Ber.  Dtsch.  chem.  Ges.,  xvi,  447. 


SEPARATION  OF  ALUMINUM  FROM  IRON.  21 

in  a  filtering  crucible  or  cone,  and  washing  the  crystalline 
precipitate  with  the  strongest  hydrochloric  acid.  Aluminum 
chloride  prepared  in  this  way  gave  no  trace  of  color  when 
dissolved  in  water  and  tested  with  potassium  sulphocyanide. 
The  correlative  question  as  to  how  much  aluminum  chloride 
goes  into  solution  under  the  conditions  was  settled  by  taking 
a  portion  of  the  pure  aluminum  chloride,  dissolving  it  in  a 
very  little  water,  diluting  the  solution  with  strong  hydrochloric 
acid,  saturating  the  cooled  liquid  with  the  gaseous  acid,  filtering 
on  asbestos,  precipitating  by  ammonia  the  aluminum  salt  in  the 
nitrate  and  weighing  the  ignited  oxide. 

From  10  cm3  of  such  a  filtrate  we  obtained  in  two  deter- 
minations 0.0022  grm.  and  0.0024  grm.  of  the  oxide,  the  mean 
of  which  corresponds  to  23  parts  of  the  oxide  or  109  parts  of 
the  hydrous  chloride  in  100,000  parts  of  the  strong  hydrochloric 
acid.  This  degree  of  solubility,  though  inconsiderable  when 
the  objective  point  is  the  preparation  of  the  pure  salt  of 
aluminum,  is  obviously  incompatible  with  the  attainment  of 
quantitative  accuracy  in  the  retention  of  the  aluminum.  We 
have  found,  however,  that  various  mixtures  of  anhydrous  ether 
and  the  strongest  hydrochloric  acid  can  be  used  satisfactorily 
as  solvents  for  the  iron  chloride,  while  the  aluminum  chloride 
is  insoluble  to  a  very  high  degree  in  a  mixture  of  hydrochloric 
acid  and  ether  taken  in  equal  parts  and  thoroughly  saturated 
with  gaseous  hydrochloric  acid  at  the  atmospheric  temperature. 
We  found  that  50  cm3  of  the  solution  of  aluminum  chloride, 
obtained  by  mixing  about  0.1  grm.  of  the  hydrous  chloride 
(dissolved  in  2  cm3  of  water)  with  the  mixture  of  pure, 
specially  prepared  aqueous  hydrochloric  acid  and  ether  in  equal 
parts  and  again  saturating  the  liquid  at  15°  C.  with  gaseous 
hydrochloric  acid,  left  upon  evaporation  and  ignition  0.0004 
grm.  in  each  of  two  experiments  —  results  which  indicate  a 
maximum  solubility  corresponding  to  1  part  of  the  oxide  or 
approximately  5  parts  of  the  chloride  in  125,000  parts  of  the 
equal  mixture  of  ether  and  aqueous  hydrochloric  acid  of  full 
strength. 

Pure  aqueous  hydrochloric   acid  of  full  strength  mixes 


22  A   METHOD  FOR   THE  SEPARATION 

perfectly  with  its  own  volume  of  anhydrous  ether,  but  it  is 
a  curious  fact  that  the  addition  to  this  mixture  of  any  very 
considerable  amounts  of  a  solution  of  ferric  chloride  in  strong 
hydrochloric  acid  determines  the  separation  of  a  greenish  oily 
ethereal  solution  of  the  ferric  salt  upon  the  surface  of  the 
acid.  The  addition  of  more  aqueous  acid  does  not  change  the 
conditions  essentially,  but  more  ether  renders  the  acid  and 
the  oily  solution  completely  miscible.  The  ferric  chloride 
seems  to  abstract  ether  from  the  ether-acid  mixture  and,  then 
dissolved  in  the  ether,  remains  to  some  extent  immiscible  with 
the  aqueous  acid  thus  left  until  the  addition  of  more  ether 
restores  to  the  mixture  that  which  was  taken  from  it  by  the 
ferric  chloride.  Our  experiments  show  that,  while  for  the 
separation  of  insoluble  aluminum  chloride  from  certain  small 
amounts  of  soluble  ferric  chloride  the  mixture  of  the  strongest 
aqueous  hydrochloric  acid  and  ether  in  equal  parts  serves  a 
most  excellent  purpose,  when  larger  amounts  of  ferric  chloride 
are  to  be  dissolved  ether  must  be  added  proportionately  in 
order  to  prevent  the  separation  of  the  ethereal  solution  of 
ferric  chloride  from  the  rest  of  the  liquid. 

Great  care  was  taken  to  insure  the  purity  of  the  aluminum 
chloride  used  in  the  test  experiments.  The  so-called  pure 
chloride  of  commerce  was  dissolved  in  the  least  possible 
amount  of  water  and  this  solution  was  treated  with  a  large 
volume  of  strong  hydrochloric  acid.  The  chloride  thus 
obtained,  free  from  iron,  but  possibly  contaminated  (as  we 
found  by  experience)  with  some  alkaline  chloride,  was  dis- 
solved in  water  and  converted  by  ammonia  to  the  form  of  the 
hydroxide,  which  was  thoroughly  washed  and  dissolved  in  hot 
hydrochloric  acid  of  half-strength.  From  this  solution,  after 
cooling,  gaseous  hydrochloric  acid  precipitated  the  hydrous 
chloride  in  pure  condition.  The  chloride  thus  prepared  was 
dissolved  in  water  and  the  strength  of  the  solution  was 
determined  by  precipitating  the  hydroxide  from  definite 
portions,  and  weighing  the  ignited  oxide  in  the  usual  manner. 

In  the  experiments  recorded  in  Table  I,  measured  portions 
of  the  standardized  solution  were  submitted  to  the  treatment 


OF  ALUMINUM  FROM  IRON. 
TABLE  I. 


23 


Bxp. 

A1203  taken 
in  solution  as 
the  chloride. 

found. 

Final 
volume. 

Error. 

grin. 

gnn. 

cm» 

grm. 

1) 

0.0761 

0.0746 

50 

0.0015- 

2 

0.0761 

0.0745 

50 

0.0016- 

3) 

0.0761 

0.0741 

50 

0.0020- 

4) 

0.0761 

0.0734 

50 

0.0027- 

5) 

0.0761 

0.0756 

50 

0.0005- 

6) 

0.0157 

0.0149 

45 

0.0008- 

7 

0.0157 

0.0147 

40 

0.0010- 

8) 

0.0157 

0.0144 

45 

0.0013- 

(9) 

0.0480 

0.0481 

30 

0.00014- 

(10) 

0.0960 

0.0957 

30 

0.0003- 

with  hydrochloric  acid  and  ether.  The  essential  thing  in  the 
process  is  to  have  at  the  end  a  mixture  of  the  strongest  aqueous 
hydrochloric  acid  with  an  equal  volume  of  anhydrous  ether 
saturated  at  a  temperature  of  about  15°  C.  The  most  con- 
venient way  to  secure  these  conditions  seems  to  be  to  mix  the 
aqueous  solution  of  the  aluminum  salt  with  a  suitable  volume 
of  the  strongest  aqueous  hydrochloric  acid  —  enough  to  make 
the  entire  volume  something  between  15  and  25  cm3  —  to 
saturate  this  mixture  with  gaseous  hydrochloric  acid  while  the 
liquid  is  kept  cool  by  immersing  the  receptacle  containing  it 
in  a  current  of  running  water,  to  intermix  a  volume  of  ether 
equal  to  the  volume  of  the  liquid,  and  finally,  to  treat  the 
ethereal  mixture  once  more  with  the  gaseous  acid  to  insure 
saturation.  The  precipitated  crystalline  chloride  was  collected 
upon  asbestos  in  a  perforated  crucible,  washed  with  a  previously 
prepared  mixture  of  hydrochloric  acid  and  ether  carefully 
saturated  with  the  gaseous  acid  at  15°  C.,  and  either  ignited 
after  careful  drying  at  150°  or  redissolved  in  water,  converted 
to  the  hydroxide  by  ammonia  in  the  usual  way  and  weighed 
as  the  oxide  after  filtration,  washing,  and  ignition.  In  experi- 
ments (1)  to  (4)  the  precipitated  chloride  was  ignited  directly ; 
in  experiment  (5)  the  ignition  was  made  with  great  care  in  an 
atmosphere  of  superheated  steam ;  and  in  experiments  (6)  to 
(10)  the  chloride  was  dissolved,  precipitated  as  the  hydroxide, 
and  weighed  as  the  oxide. 


24  A   METHOD  FOR   THE  SEPARATION 

The  experiments  in  which  the  chloride  was  converted  to  the 
hydroxide  before  ignition  show  upon  the  average  an  absolute 
loss  of  about  0.0006  grm. :  the  single  experiment  in  which  the 
ignition  took  place  in  steam  shows  about  the  same  loss  —  0.0005 
grm. ;  while  in  those  experiments  in  which  the  chloride  was 
dried  and  then  ignited  directly,  the  average  loss  amounts  to 
about  0.0020  grm.  The  error  of  the  process  which  involves 
the  precipitation  of  the  aluminum  as  the  hydroxide,  falls  within 
reasonable  limits,  but  it  is  plain  that  the  direct  ignition  of  the 
chloride  is  liable  to  error,  which  may  possibly  be  explicable  as 
a  mechanical  loss  occasioned  by  the  too  rapid  evolution  of  the 
hydrochloric  acid  and  water  of  crystallization,  or,  possibly,  as 
the  result  of  a  very  slight  volatilization  of  the  aluminum  still 
holding  chlorine  in  spite  of  the  decomposing  action  of  the 
water  upon  the  chloride.  In  either  case,  it  would  seem  to  be 
reasonable  to  suppose  that  a  layer  of  some  easily  volatilizable 
oxidizer  placed  upon  the  aluminum  chloride  might  serve  to 
obviate  the  difficulty  —  in  the  one  case,  by  serving  as  a  screen 
to  diminish  mechanical  transportation  of  the  non-volatile 
material ;  and  in  the  other,  by  acting  as  an  agent  to  promote 
the  exchange  of  chlorine  for  oxygen  on  the  part  of  the 
aluminum  chloride. 

We  have  tried,  therefore,  the  expedient  of  covering  the 
aluminum  chloride  before  ignition  with  a  layer  of  mercuric 
oxide,  which  of  itself  left  no  appreciable  residue  when  it 
volatilized.  The  hydrous  chloride  was  collected  as  usual  upon 
the  asbestos  in  a  perforated  crucible,  dried  for  a  half-hour  at 
150°  C.,  covered  with  about  1  grm.  of  the  pure  mercuric  oxide, 
gently  heated  with  great  care  under  a  suitable  ventilating  flue, 
and  finally  ignited  over  the  blast.  The  results  are  given 
below  (see  Table  II). 

It  is  obvious,  therefore,  that  the  precipitation  of  the  crystal- 
line hydrous  aluminum  chloride  from  solutions  of  the  pure  salt 
is  perfectly  feasible  and  very  complete,  when  effected  by 
aqueous  hydrochloric  acid  and  ether  thoroughly  saturated  with 
the  gaseous  acid  and  kept  cool ;  and  that  the  conversion  of  the 
chloride  into  the  weighable  form  of  the  oxide  is  best  effected 


OF  ALUMINUM  FROM  IRON. 


25 


TABLE   II. 


Exp. 

Al,0,,  taken 
in  solution  as 
the  chloride. 

AlaO3  found 
by  ignition 
with  HgO. 

Final 
volume. 

Error. 

grm. 

grm. 

cm3 

grm. 

(1) 

0.0761 

0.0758 

25 

0.0003- 

0.0761 

0.0754 

25 

0.0007- 

(3) 

0.0761 

0.0761 

25 

0.0010- 

by  ignition  under  a  layer  of  mercuric  oxide,  or  by  dissolving 
it  in  water  and  precipitating  it  as  the  hydroxide  to  be 
afterward  washed,  dried,  and  ignited.  Of  the  two  methods  the 
former  is  by  far  the  more  convenient. 

The  precipitation  of  the  aluminum  chloride  in  pure  condition 
from  solutions  containing  ferric  chloride  ought  not,  it  would 
seem,  to  present  any  difficulty,  providing  only  that  the 
precaution  is  taken  to  have  present  a  sufficient  excess  of  ether. 
The  question  was  put  to  the  test  of  experiment  with  the 
results  recorded  in  Table  III. 

TABLE  III. 


Exp. 

A1,O3  taken 
in  solution  as 
the  chloride. 

A15O3  found 
by  ignition 
with  HgO. 

Fe203 
present  as 
chloride. 

Final 
volume. 

Error. 

grin. 

grm. 

grm. 

cm3 

grm. 

(1) 

0.0761 

0.0757 

0.15 

25-30 

0.0004- 

2 

0.0761 

0.0756 

0.15 

25-30 

0.0005- 

3 

0.0761 

0.0755 

0.15 

25-30 

0.0006- 

4) 

0.0761 

0.0755 

0.15 

25-30 

0.0006- 

Measured  portions  of  the  standardized  solution  of  aluminum 
chloride  were  evaporated  nearly  to  dryness  in  a  platinum  dish, 
an  amount  of  pure  ferric  chloride  equivalent  to  about  0.15 
grm.  of  the  oxide  was  added  in  a  very  little  water,  15  cm3  of 
the  mixture  of  strong  hydrochloric  acid  and  ether  in  equal 
parts  were  introduced,  the  liquid  was  saturated  at  15°  C.  with 
gaseous  hydrochloric  acid  (the  dish  being  held  in  a  convenient 
device  for  cooling  it  by  running  water),  5  cm3  more  of  ether 
were  added  to  secure  complete  miscibility  of  the  solutions,  and 


26  SEPARATION  OF  ALUMINUM  FROM  IRON. 

more  gas  passed  to  perfect  saturation.  The  aluminum  chloride 
was  collected  upon  asbestos  in  a  perforated  crucible,  washed 
with  a  mixture  of  ether  and  aqueous  hydrochloric  acid 
thoroughly  saturated  with  the  gaseous  acid,  dried  at  150°  C. 
for  a  half -hour,  covered  with  1  grm.  of  pure  mercuric  oxide, 
and  ignited  at  first  gently  and  finally  over  the  blast. 

The  results  show  plainly  a  very  satisfactory  limit  of  error. 


THE  ESTIMATION  OF  MOLYBDENUM 
IODOMETRICALLY. 

BY  F.  A.  GOOCH  * 

IN  a  former  paper  from  this  laboratory  f  several  modes  of 
applying  hydriodic  acid  to  the  reduction  of  molybdic  acid 
were  studied.  It  was  found,  first,  that  the  digestion  process 
of  Mauro  and  Danesi  f  is  of  very  limited  applicability, 
owing  to  the  fact  that  the  reaction  of  reduction  is  reversible. 
Secondly,  it  appeared  that  the  use  of  the  same  reaction  by 
Friedheim  and  Euler  §  in  a  distillation  process,  so  arranged 
that  the  iodine  set  free  in  the  reduction  might  be  caught  in 
the  distillate  and  titrated  to  serve  as  the  measure  of  the 
reducing  action,  was  not  sufficiently  regular  because  of  in- 
attention to  minor  details.  It  was  shown  that  by  taking 
care  to  adjust  the  conditions  constant  results  might  be 
obtained.  Thirdly,  the  fact  was  developed  that  by  simply 
boiling  the  solution  under  well  defined  conditions  in  an 
ordinary  Erlenmeyer  flask,  partly  closed  by  a  simple  trap, 
the  reduction  of  the  molybdic  acid  proceeded  regularly,  and 
that  the  addition  of  standard  iodine  to  the  solution  made 
alkaline  with  sodium  bicarbonate  served  to  restore  the  original 
condition  of  oxidation  of  the  molybdic  acid.  The  results 
of  this  treatment  were  shown  to  be  accurate. 

In  a  recent  paper  ||  Friedheim  has  seen  fit  to  make  our 
modifications  of  the  distillation  process  the  subject  of  attack. 
Friedheim's  comments  upon  the  third  method  discussed  (as 

*  From  Am.  Jour.  Sci.,  iii,  237. 

t  Gooch  and  Fairbanks,  Am.  Jour.  Sci.,  ii,  157.    Volume  I,  p.  375. 

J  Zeitschr.  anal.  Chem..  xx,  507. 

§  Ber.  Dtsch.  chem.  Ges.,  xxviii,  2066. 

||  Ber.  Dtsch.  chem.  Ges.,  xxix,  2981. 


28  THE  ESTIMATION  OF  MOLYBDENUM. 

well  as  upon  a  subsequent  application  of  the  process)  *  are 
evidently  prompted  wholly  by  personal  opinion  and  demand 
no  further  attention.  With  reference  to  Friedheim's  denial 
of  the  necessity  of  modification  in  the  Friedheim  and  Euler 
treatment  the  case  is  different. 

The  process  of  Friedheim  and  Euler  consists,  it  will  be 
remembered,  in  treating  the  soluble  molybdate,  or  the  solu- 
tion of  molybdic  acid  in  sodium  hydroxide,  with  potassium 
iodide  and  hydrochloric  acid  in  a  Bunsen  apparatus,  boiling 
until  the  solution  is  of  a  clear  green  color,  collecting  the 
iodine  distilled  in  potassium  iodide,  and  titrating  it  with 
sodium  thiosulphate.  We  found  that  the  development  of 
the  green  color  was  not  a  sufficient  criterion  of  the  exact 
reduction  of  the  molybdic  acid  to  the  condition  of  the 
pentoxide  and  of  the  removal  of  the  iodine  which  should  be 
theoretically  set  free.  To  accomplish  that  end  we  found  it 
safer  and  more  convenient  to  start  the  distillation  with  a 
definite  volume  (40  cm3)  of  liquid  and  boil  until  a  definite 
volume  (25  cm3)  was  reached,  care  being  taken  with  regard 
to  the  strength  of  acid  and  the  excess  of  potassium  iodide 
employed.  Experience  showed  unmistakably  that  in  order 
to  avoid  the  decomposing  action  of  the  air  upon  the  hot 
vaporous  hydriodic  acid  in  the  retort,  it  was  necessary  to 
go  beyond  the  measures  advised  by  Friedheim  and  Euler 
(namely,  to  warm  the  retort  and  its  contents  slowly,  heating 
to  boiling  only  when  the  connecting  tube  was  well  filled 
with  iodine  vapor  and  the  tendency  toward  back-suction  of 
the  liquid  in  the  receiver  began  to  appear)  and  to  conduct 
the  operation  in  a  simple  little  apparatus  (the  retort  holding 
about  100  cm3)  put  together  entirely  with  sealed  and  ground 
joints,  as  shown  in  the  figure  of  the  former  paper,  so  ar- 
ranged that  a  current  of  purified  carbon  dioxide  could  be 
passed  through  retort  and  receiver  during  the  distillation. 
With  this  apparatus  we  were  able  to  determine  with  accuracy 
the  point  of  concentration  at  which  the  free  iodine  left  the 
liquid,  the  molybdic  acid  having  been  converted  to  the  con- 
*  Am.  Jour.  jSci.  ii,  181.  Volume  I,  p.  391. 


THE  ESTIMATION  OF  MOLYBDENUM.  29 

dition  of  the  pentoxide.  It  was  found  that  if  dependence  is 
placed  upon  the  occurrence  of  the  so-called  clear  green  color 
of  the  liquid  to  determine  the  end  of  the  distillation,  it  may 
frequently  happen  that  free  iodine  remains  in  the  residue. 
This  takes  place,  it  will  be  observed,  in  the  atmosphere  of 
carbon  dioxide,  so  that  the  presence  of  the  free  iodine  can 
by  no  possibility  be  attributed  to  the  action  of  atmospheric 
air  upon  the  hydriodic  acid  remaining  after  the  distillation 
is  complete.  On  the  other  hand,  it  appeared  that,  if  the 
distillation  is  pushed  too  far,  the  molybdenum  pentoxide  may 
be  still  further  reduced  with  consequent  evolution  of  more 
than  the  expected  amount  of  iodine.  The  attainment  of  an 
exact  degree  of  reduction  with  the  expulsion  of  the  corre- 
sponding amount  of  iodine  becomes,  therefore,  a  matter  of 
chance  unless  further  precautions  are  taken.  We  found  in 
our  experiments  that,  if  amounts  less  than  0.3  grm.  of  the 
molybdic  acid  are  introduced  in  soluble  form  into  the  100  cm3 
retort  with  a  not  too  great  excess  of  potassium  iodide,  and 
the  40  cm3  of  liquid  so  constituted  that  20  cm3  of  it  shall  be 
water  and  20  cm  3  the  strongest  hydrochloric  acid,  the  reduc- 
tion proceeds  with  a  fair  degree  of  regularity  in  the  manner 
expected.  We  found  it  important  to  restrict  the  excess  of 
potassium  iodide  so  that  it  shall  never  exceed  the  theoretical 
requirement  by  more  than  0.5  grm. 

Our  determinations  with  the  pure  molybdenum  trioxide 
showed  errors  varying  from  0.0010  grm.  -|-  to  0.0007  grm.  — ; 
the  variations  from  theory  in  the  experiments  with  ammonium 
molybdate  ranged  from  0.0011  grm.  -f-  to  0.0011  grm.  — .  If 
these  results  are  compared  with  those  given  by  Friedheim 
and  Euler,  the  advantage  is  a  little  in  favor  of  the  latter; 
but  a  scrutiny  of  the  figures  given  by  Friedheim  and  Euler 
develops  the  fact  that  the  apparent  accuracy  of  their  work 
is  founded  upon  miscalculations.  This  fact  was  known  to  us 
at  the  tune  of  our  former  writing,  but  we  did  not  consider  it 
essential  then  to  make  the  matter  public.  The  recent  attack 
of  Friedheim  makes  that  course  now  necessary. 

Herewith  is  reproduced  a  table  of  results  obtained  by  Fried- 


30 


THE  ESTIMATION  OF  MOLYBDENUM. 


heim  and  Euler  in  the  test  of  their  method  upon  ammonium 
molybdate,  shown  by  analysis  to  contain  81.49  per  cent  of 
molybdenum  trioxide.  The  figures  which  are  incorrect  are 
enclosed  in  brackets: 

OEIGINAL  FIGUKES  OF  FRIEDHEIM  AND  EULEB. 


Per  cent  of 

Molybdate 
taken. 

Na2S203 
used. 

Mo08 
found. 

MoOs  referred 
to  molybdate 

taken. 

grm. 

cm8 

grm. 

0.2674 

30.8     )    1  cm2  = 

0.2184 

[81.71] 

0.4418 

50.8     >    0.00709 

0.3601 

81.51 

0.4075 

[40.71*1    Mo08. 

0.3317 

81.40 

0.3281 
0.4340 
0.4098 
0.4305 

37.33  1    T       o 

ACk  AQ      I       1   Cm     = 

I?'!?  [   0.007086 
46.63   I    ii/r^r* 
49.08  J    Mo°«- 

0.2644 
0.3502 
0.3304 
0.3478 

t  81.85-1 
81.69 
81.67 
81.78J 

Appended  is  a  recalculation  of  the  percentage  of  the  trioxide 
found,  with  columns  showing  the  percentage  error  and  the 
error  stated  in  fractions  of  a  gram.  Changes  from  the  figures 
of  Friedheim  and  Euler  are  in  heavy-faced  type. 

RECALCULATION  OF  THE  RESULTS  OF  FRIEDHEIM  AND  EULEB. 


Corrected 
per  cent  of  MoO», 
found,  referred  to 
the  molybdate. 

Error  in 
per  cent  of 
Mo03  found 
compared  with 
Mo03  taken. 

Error 
of  MoOa. 

grm. 

81.68 

0.23+ 

0.0005+ 

81.51 

0.03+ 

0.0001+ 

81.40 

0.12- 

0.0004- 

80.58 

1.12- 

0.0030- 

80.69 

0.99- 

0.0035- 

80.62 

1.05- 

0.0035- 

80.79 

0.86- 

0.0030- 

These  figures  of  their  own  (properly  calculated)  are  suffi- 
cient to  show  the  inadequacy  of  the  method  of  Friedheim  and 
Euler.  We  ourselves  were  occasionally  able  to  get  results 
from  the  method  of  Friedheim  and  Euler  quite  as  good  as 


Probably  46.7. 


THE  ESTIMATION  OF  MOLYBDENUM. 


31 


these;  it  must  be  said,  however,  that  most  of  our  results 
obtained  by  their  unmodified  method  have  been  even  worse 
than  their  own. 

In  another  series  of  six  determinations,  in  which  molybde- 
num trioxide  was  the  starting-point,  Friedheim  and  Euler 
were  more  successful,  the  errors  varying  from  0.0006  grm.  + 
to  0.0006  grm.—.  Thus  Friedheim  and  Euler  establish  by 
their  own  results  the  fact  that  the  hitting  of  the  right  point  at 
which  to  stop  their  process  of  boiling  is  a  matter  of  chance. 
In  spite  of  the  probability  that  some  of  the  iodine  which  they 
found  in  the  receiver  was  liberated  by  atmospheric  action,  the 
fact  remains  that  their  results  are  in  many  cases  very  low. 
That  is,  they  did  not  boil  long  enough. 

The  difficulty  appears  again  in  the  modification  of  their 
method  which  Friedheim  and  Euler  apply  to  the  determination 
of  molybdenum  trioxide  associated  with  vanadium  pentoxide,* 
namely,  the  distillation  with  phosphoric  acid  and  potassium 
iodide  of  the  residue  left  after  reducing  the  vanadium  pentox- 
ide by  hydrochloric  acid  and  potassium  bromide,  according  to 
the  method  of  Holverscheit.  We  reproduce  the  part  of  their 
table  which  refers  to  the  determination  of  the  molybdenum, 
adding,  however,  columns  containing  the  errors  and  corrected 
percentages. 


MoOa 
taken. 

MoO? 

Per  cent 
Mo03 

Error. 

Per  cent 
Mo03. 

P.  and  E. 

Recalculated. 

grin. 

gnu. 

grm. 

0.15037 

0.15005 

99.79 

0.00032- 

99.79 

0.16895 

0.16879 

99.90 

0.00016- 

99.90 

0.17758 

0.17729 

99.84 

0.00029— 

99.84 

0.24975 

0.24962 

99.95 

0.00013- 

99.95 

0.33151 

0.33607 

[99.87] 

0.00456+ 

101.38 

Four  of  the  five  determinations  are  accurate,  but  the  fact 
that  all  figures  are  carried  out  to  the  fifth  decimal  place  does 
not  keep  three  good-sized  figures  out  of  the  error  column  for 
the  fifth  determination. 


*  Ber.  Dtsch.  chem.  Ges.,  xxviii.  2072. 


32  THE  ESTIMATION  OF  MOLYBDENUM. 

It  is  hardly  necessary,  in  the  light  of  a  comparison  of  the 
results  of  Friedheim  and  Euler  with  ours,  to  discuss  further 
the  unreliability  of  the  unmodified  process.  The  necessity  of 
a  proper  control  of  the  volume,  strength  of  acid,  and  excess 
of  potassium  iodide,  as  well  as  proper  protection  from  atmos- 
pheric oxidation,  is  real. 


VI 


THE  APPLICATION   OF  IODIC   ACID  TO   THE 
ANALYSIS   OF   IODIDES. 

BY  F.  A.  GOOCH  AND  C.  F.  WALKER.* 

IT  has  long  been  understood  that  iodic  acid  is  easily  and 
completely  reduced  by  an  excess  of  hydriodic  acid  with  the 
liberation  of  iodine  according  to  the  equation: 

HI08  +  5HI  =  61  +  3H2O. 

To  apply  this  reaction  to  the  quantitative  estimation  of  iodic 
acid,  it  is  only  necessary  to  add  to  the  free  iodic  acid  or  solu- 
ble iodate  an  excess  of  a  soluble  iodide,  to  acidify  —  best  with 
dilute  sulphuric  acid  —  and  to  titrate  the  iodine  thus  set  free 
with  sodium  thiosulphate,  one-sixth  of  the  iodine  found  being 
credited  to  the  iodic  acid. 

It  has  been  shown  recently  by  Rieglerf  that  this  reaction 
may  be  also  applied  to  the  quantitative  estimation  of  iodides, 
the  iodine  set  free  upon  the  addition  of  a  known  excess  of 
iodic  acid  to  the  iodide  solution  being  removed  by  petroleum 
ether,  and  the  residual  iodic  acid  titrated  directly  with  sodium 
thiosulphate. 

The  present  investigation  was  undertaken  to  define  more 
particularly  the  limit  of  applicability  of  the  reaction  and  to 
establish,  if  possible,  a  direct  method  for  the  quantitative  esti- 
mation of  iodides,  dependent  upon  the  action  of  iodic  acid  or 
an  iodate  in  the  presence  of  free  sulphuric  acid,  neutralization 
of  the  solution  by  means  of  an  acid  carbonate,  and  titration  of 
the  free  iodine  by  arsenious  acid  —  five-sixths  of  the  iodine 
thus  found  being  credited  to  the  iodide  to  be  estimated.  It 

*  From  Am.  Jour.  Sci.  iii,  293.  t  Zeitschr.  anal.  Chem.,  xxxv,  305. 

VOL.    II.  —  3 


34  THE  APPLICATION  OF  10DIC  ACID 

has  been  found  that  by  fulfilling  certain  necessary  conditions, 
the  proposed  method  is  entirely  successful,  so  far  as  concerns 
the  estimation  of  iodine  in  iodide  solutions  free  from  large 
amounts  of  chlorides  or  bromides. 

In  a  system  containing  a  considerable  quantity  of  free 
iodine  with  variable  amounts  of  the  other  reagents  mentioned, 
as  well  as  possible  impurities,  it  is  conceivable  that  secondary 
reactions  may  occur,  depending  largely  on  conditions  of  mass, 
tune,  and  temperature,  and  of  a  sort  likely  to  alter  the  amount 
of  recoverable  iodine,  or  to  exert  an  excessive  oxidizing  influ- 
ence on  the  arsenious  acid  finally  titrated.  It  has  been  estab- 
lished by  Schonbein,  Lunge  and  Schloch,  and  others,  that 
iodine  forms  compounds  with  the  alkalies  of  the  type  R-O-I, 
and  Phelps*  has  recently  found  that  the  formation  of  some 
such  compound,  accompanying  the  iodate  naturally  expected, 
is  distinctly  recognizable  when  iodine  and  barium  hydroxide 
interact  at  ordinary  temperatures.  It  has  been  shown,  also, 
in  a  former  paper  from  this  laboratory!  that  free  iodine  or  an 
iodide  interacts  very  easily  with  iodic  acid  in  the  presence  of 
dilute  hydrochloric  acid  with  the  formation  of  iodine  mono- 
chloride,  according  to  the  equations: 

HI03  +  21,  +  5HC1  =  3H20  +  6IC1. 

HI03  +  2KI  +  5HC1  =  3H20  +  2KC1  +  3IC1. 

Moreover,  organic  compounds  containing  the  groups  —1  =  0 
and  —I  ~  Q,  in  which  iodine  seems  to  be  analogous  to  nitrogen, 

result  in  great  variety  from  the  oxidation  of  halogen  sub- 
stitution products.  It  would  seem,  therefore,  that  the 
formation  of  inorganic  reduction  products  of  iodic  acid  under 
the  conditions  likely  to  obtain  in  this  analytical  process  might 
be  by  no  means  beyond  the  bounds  of  possibility. 

A  few  simple  qualitative  tests  to  determine  the  possibility 
of  interaction  between  small  quantities  of  iodine  and  iodic 
acid  alone  met  with  negative  results.  Thus,  a  single  drop 
of  a  decinormal  solution  of  iodine,  made  as  usual  in  potassium 

*  Am.  Jour.  Sci.,  ii,  70.     Volume  I,  p.  370. 

t  Roberts,  Am.  Jour.  Sci.,  xlviii,  157.    Volume  I,  p.  257. 


TO   THE  ANALYSIS  OF  IODIDES.  35 

iodide,  gave  when  added  to  10  cm3  of  decinormal  iodic  acid  a 
distinctive  color  to  chloroform.  Similar  results  were  obtained 
when  the  iodine  was  employed  in  aqueous  solution  in  which 
there  was  no  alkaline  iodide.  A  few  drops  of  an  aqueous 
solution  of  iodine  treated  (in  either  order)  with  10  cm3  of  a 
saturated  solution  of  potassium  bicarbonate  and  10  cm3  of 
decinormal  iodic  acid  gave  the  same  distinctive  color  to 
chloroform  as  came  from  the  same  amount  of  iodine  in  the 
absence  of  the  iodic  acid.  So  it  appears  that  if  in  the  system 
under  consideration  reactions  do  occur  between  iodic  acid  and 
iodine  to  alter  the  amount  of  iodine  recoverable,  such  action 
is  not  appreciable  between  small  amounts  of  these  materials. 
This,  however,  does  not  preclude  the  possibility  of  perceptible 
changes  under  the  mass-action  of  a  large  amount  of  iodine. 

The  reactions  of  hydrochloric  acid,  and  probably  of  hydro- 
bromic  acid,  hi  the  presence  of  varying  amounts  of  iodic 
acid,  iodine,  and  iodide,  as  well  as  the  reaction  of  the  alkaline 
carbonate  upon  such  mixtures  are  doubtless  complex,  more  or 
less  reversible,  and  dependent  upon  proportions  and  dilution. 
The  tendency  of  the  former  reactions  is  toward  the  reduction 
of  the  molecule  of  iodic  acid,  and  the  formation  of  the  chloride 
or  bromide  of  iodine.  Thus,  Miss  Roberts  *  demonstrated 
that  a  solution  of  hydrochloric  acid,  so  dilute  that  by  itself  it 
is  without  effect  on  iodic  acid,  acts  upon  a  mixture  of  iodic 
acid  with  either  free  iodine  or  an  iodide  to  form  iodine 
monochloride.  The  action  of  the  acid  carbonate  upon  the 
iodine  chloride  or  bromide  may  produce  a  salt  of  the  oxy-acids 
and  free  iodine. 

The  practical  effects,  under  the  conditions  of  analysis,  of  the 
reaction  between  iodine,  iodic  acid  and  the  halogen  acids  in 
presence  of  sulphuric  acid,  and  of  reactions  which  may  occur 
upon  neutralization  by  an  acid  carbonate,  were  studied  in  detail 
in  a  number  of  experiments. 

The  preliminary  experiments  of  Table  I  were  made  to  bring 
out  the  effect  of  neutralizing  with  the  acid  carbonate  and 
subsequently  titrating  with  an  alkaline  arsenite  a  solution 

*  Loc.  cit. 


36 


THE  APPLICATION  OF  IODIC  ACID 


[5  cm3 


TABLE  I. 

EFFECT  OF  THE  CARBONATE. 
(1:3).    Total  volume  of  liquid,  250  cm8.] 


Ezp. 

I  (in  KI)  taken. 

KHC03  in  excess. 

I  found. 

Error. 

grin. 

cm8 

grin. 

gnu. 

(1) 

0.0713 

Very  small. 

0.0707 

0.0006- 

(2) 

0.0715 

Very  small. 

0.0710 

0.0005- 

(3) 

0.0713 

10 

0.0710 

0.0003- 

(4) 

0.0710 

10 

0.0706 

0.0004- 

(5) 

0.0723 

10 

0.0717 

0.0006- 

(6) 

0.0713 

20 

0.0709 

0.0004- 

(7) 

0.0713 

20 

0.0709 

0.0004- 

(8) 

0.3565 

Very  small. 

0.3560 

0.0005- 

(9) 

0.3568 

Very  small. 

0.3561 

0.0007- 

(10) 

0.3667 

10 

0.3563 

0.0004- 

(11) 

0.3596 

10 

0.3588 

0.0008- 

(12) 

0.3565 

10 

0.3565 

0.0000 

(13) 

0.3572 

20 

0.3560 

0.0012- 

(14) 

0.3567 

20 

0.3569 

0.0002-f- 

containing  sulphuric  acid  and  a  considerable  amount  of  free 
iodine.  The  danger  of  mechanical  loss  of  iodine  during  the 
effervescence  accompanying  neutralization,  as  well  as  by 
spontaneous  volatilization  from  the  surface  during  the  process 
of  titration,  was  minimized  by  effecting  the  neutralization  in 
the  trapped  Drexel  washing-bottle  to  be  described  later,  and 
making  the  titration  in  the  same  tall  washing  cylinder  without 
transfer.  To  varying  amounts  of  a  recently  standardized 
decinormal  solution  of  iodine  were  added  successively  5  cm3 
of  dilute  sulphuric  acid  and  varying  amounts  of  potassium 
bicarbonate  in  excess  of  that  necessary  to  neutralize  the  free 
acid,  decinormal  arsenious  acid  in  slight  excess  of  the  iodine, 
5  cm3  of  starch  emulsion,  and  decinormal  iodine  to  coloration, 
the  total  volume  of  the  liquid  being  not  greater  than  250  cm8. 
The  results  show  plainly  that  while  the  loss,  mechanical  or 
otherwise,  in  the  treatment  of  reasonably  large  amounts  of 
fairly  concentrated  iodine  is  perceptible,  it  is  still  well  within 
permissible  limits  (amounting  to  a  little  less  than  0.0005  grm. 
in  the  mean),  and  obviously  independent  of  the  excess  of  the 
carbonate  in  the  solution,  and  of  the  amount  of  free  iodine 
present. 


TO   THE  ANALYSIS  OF  IODIDES. 


37 


TABLE  IL 
EFFECT  OF  DILUTION. 


Bxp. 

El  taken. 

KI  found. 

Error. 

Approximate 
volume  upon 
addition  of  H2S04. 

Volume 
H2S04(1:3) 
used. 

(1) 

(2) 
(3) 
(4) 
(5) 
(6) 
(7) 
(8) 

grin. 

0.0772 
0.0772 
0.1544 
0.1544 

0.3087 
0.3087 
0.3859 
0.3859 

grlu. 

0.0768 
0.0765 
0.1546 
0.1541 
0.3090 
0.3088 
0.3864 
0.3860 

grm. 
0.0004- 
0.0007- 
0.0002+ 
0.0003- 
0.0003+ 
0.0001+ 
0.0005+ 
0.0001+ 

1SSSSSSSS 

cm** 

6 
5 
5 
5 
5 
6 
5 
5 

(9) 
(10) 

(11) 
(12) 

0.0772 
0.0772 
0.1543 
0.1544 

0.0754 
0.0757 
0.1532 
0.1524 

0.0018- 
0.0015- 
0.0011- 
0.0020- 

300 
300 
300 
300 

5 
5 
5 
5 

(13) 
(14) 
(16) 
(16) 
(17) 
(18) 

0.0772 
0.0772 
0.1544 
0.1544 
0.3859 
0.3859 

0.0744 
0.0737 
0.1521 
0.1512 

0.3827 
0.3831 

0.0028- 
0.0035- 
0.0023- 
0.0032- 
0.0032- 
0.0028- 

600 
500 
500 
500 
500 
500 

5 
6 
5 
5 
5 
5 

(19) 
(20) 
(21) 
(22) 

0.0772 
0.0772 
0.3859 
0.3859 

0.0744 
0.0757 

0.3828 
0.3827 

0.0028- 
0.0015- 
0.0031- 
0.0032- 

500 
500 
500 
500 

10 
10 
10 

10 

In  the  experiments  of  Table  II  the  proposed  process  of 
analysis  was  tested  upon  potassium  iodide  taken  by  itself  in 
varying  amounts  of  a  J^  normal  solution  and 
carefully  standardized  by  the  method  formerly 
elaborated  in  this  laboratory.*  The  apparatus 
employed  was  a  Drexel  washing-bottle  of  500 
cm3  or  1000  cm3  capacity,  according  to  require- 
ments, with  stop-cock  and  thistle-tube  fused  to 
the  inlet  tube  and  a  Will  and  Varrentrapp 
absorption  trap  sealed  to  the  outlet,  as  shown 
in  the  accompanying  figure.  The  iodide  for  the 
test  was  drawn  from  a  burette  into  the  bottle 
and  carefully  washed  down,  and  potassium 
iodate  in  excess  of  the  amount  theoretically  necessary  (namely, 


FIG.  18. 


*  Gooch  and  Browning,  Am.  Jour.  Sci.,  xxxix,  188.    Volume  I,  p.  1. 


38  THE  APPLICATION  OF  IODIC  ACID 

5  cm3  of  a  0.5  per  cent  solution  for  every  portion  of  20  cm3 
of  the  iodide  solution),  was  added  and  the  volume  of  the 
liquid  was  adjusted  to  the  volume  at  which  it  was  desired 
that  the  iodic  and  hydriodic  acids  should  react.  The  stop- 
per with  the  thistle-tube  and  trap  was  now  placed  on  the 
bottle  and  the  trap  was  half  filled  by  means  of  a  pipette 
with  a  5  per  cent  solution  of  potassium  iodide.  Five  cen- 
timeters of  dilute  (1 : 3)  sulphuric  acid  were  added  through 
the  thistle-tube  and  washed  down ;  the  stop-cock  was  closed, 
.and  the  solution  gently  agitated,  if  necessary,  to  insure  a 
;  complete  separation  of  iodine.  Potassium  bicarbonate  in 
saturated  solution  to  an  amount  about  10  cm8  in  excess  of 
that  required  to  neutralize  5  cm8  of  dilute  (1 : 3)  sulphuric 
acid,  was  poured  into  the  thistle-tube,  and  allowed  to  flow 
into  the  bottle  slowly  enough  to  avoid  a  too  violent  evolution 
of  gas.  The  stop-cock  was  closed  and  the  solution  agitated 
by  giving  to  the  bottle  a  rotary  motion,  at  the  same  time 
keeping  the  bottom  pressed  down  upon  the  work-table,  to 
prevent  a  possible  splashing  of  the  iodide  out  of  the  trap  into 
the  yet  acid  solution.  When  the  neutralization  of  the  solution 
had  been  completed,  the  bottle  was  shaken  until  the  last  trace 
of  violet  vapor  was  absorbed  in  the  liquid.  The  greater  part 
of  the  solution  in  the  trap  was  then  run  back  into  the  bottle, 
the  stopper  removed,  and  the  tube  and  trap  carefully  washed, 
the  washings  being  added  to  the  bulk  of  the  solution.  Deci- 
normal  arsenious  acid  was  introduced  from  a  burette  to  the 
bleaching  point,  5  cm3  of  starch  emulsion  were  added,  and  the 
solution  was  titrated  back  with  decinormal  iodine  (usually 
only  a  few  drops)  to  coloration. 

Blank  tests  made  upon  a  solution  obtained  by  mixing  the 
maximum  amount  of  the  iodate  with  5  cm3  of  dilute  sulphuric 
acid  (1  : 3),  neutralizing  as  usual  with  potassium  bicarbonate, 
adding  the  iodide  from  the  trap  and  5  cm3  of  starch  emulsion, 
showed  that  a  single  drop  of  iodine  was  invariably  sufficient 
to  bring  out  the  starch  blue.  Occasionally  it  was  found  that 
the  mixture,  particularly  when  chlorides  or  bromides  were 
present,  of  itself  developed  a  trace  of  color,  but  by  no  means 


TO  THE  ANALYSIS   OF  IODIDES. 


39 


a  reading  tint.  A  correction  of  the  one  drop  of  iodine 
necessary  to  bring  out  the  color  reaction  in  the  blanks,  was 
applied  uniformly  in  the  analytical  process. 

The  number  of  centimeters  of  decinormal  arsenious  acid 
required  to  bleach  the  free  iodine,  multiplied  by  0.01383 
(log.  2.140822)  gives  the  number  of  grams  of  potassium 
iodide  taken  for  analysis,  being  equivalent  to  five-sixths  of 
the  iodine  liberated  in  the  solution. 

From  these  results  it  appears  that  the  degree  of  dilution 
of  the  solution  at  the  time  when  the  mixed  iodide  and  iodate 
are  acidified  has  an  important  influence  on  the  completeness 
of  the  reaction.  Thus,  the  mean  error  of  the  determinations 
in  which  the  volume  at  the  time  of  the  reaction  did  not 
exceed  150  cm3  was  practically  nothing,  while  the  errors  at 
volumes  of  300  cm3  and  500  cm3  amounted  to  0.0016  grm.  and 
0.0028  grm.  respectively.  It  is  obvious  that  the  doubling 
of  the  amount  of  sulphuric  acid  used  in  acidifying  does  not 
increase  the  amount  of  iodine  liberated  at  the  highest  dilution. 
The  plain  inference  is  that  the  interaction  between  the  iodide 
and  iodate  should  be  brought  about  in  a  volume  of  liquid 
not  much  exceeding  150  cm3. 

In  the  following  series  of  experiments,  recorded  in  Table 
III,  the  effect  of  the  introduction  of  a  chloride  or  bromide 

TABLE  III. 
EFFECT  OF  CHLORIDE  AND  BROMIDE. 


Exp 

KI  taken. 

KI  found. 

Error. 

NaCl  taken. 

KBr  taken. 

grin. 

grm. 

grm. 

grm. 

grm. 

(1) 

0.0772 

0.0795 

0.0023+ 

0.2 

, 

(2 

0.0772 

0.0784 

0.0012+ 

0.2 

(3 
(4 

0.0771 
0.0773 

0.0823 
0.0819 

0.0052+ 
0.0046+ 

0.5 
0.6 

(6 

0.1544 

0.1588 

0.0044+ 

0.5 

( 

6 

0.1544 

0.1590 

0.0046+ 

0.5 

o 

0.0772 

0.0802 

0.0030+ 

,  , 

6.2 

(8 

0.0773 

0.0853 

0.0080+ 

.  , 

0.2 

(9) 

0.0772 

0.0873 

0.0101+ 

t 

0.5 

(10) 

0.0772 

0.0861 

0.0089+ 

0.5 

11) 

0.1544 

0.1646 

0.0102+ 

t 

0.5 

(12) 

0.1543 

0.1626 

0.0083+ 

•  • 

0.6 

40 


THE  APPLICATION  OF  IODIC  ACID 


into  the  iodide  (before  the  iodate  is  added)  was  studied.  The 
volume  of  the  liquid  at  the  time  of  acidifying  was  fixed  at 
150  cm3  approximately,  and  5  cm3  of  the  dilute  sulphuric  acid 
(1  : 3)  were  used.  The  mode  of  procedure  was  otherwise 
similar  to  that  of  the  foregoing  series. 

The  influence  of  sodium  chloride  and  potassium  bromide 
in  increasing  the  amount  of  iodine  liberated  is  plain.  The 
increase  comes  without  doubt  from  the  iodate,  and  is  doubt- 
less due  to  the  formation  of  iodine  chloride  or  bromide,  during 
the  acidifying,  by  the  interaction  of  the  free  iodine,  the  iodic 
acid,  and  the  hydrochloric  or  hydrobromic  acid,  according  to 
the  reactions  previously  discussed.  It  is  plain,  therefore, 
that  the  value  of  the  process  in  the  determination  of  iodine 
in  an  iodide  is  restricted  of  necessity  to  those  cases  in  which 
it  is  known  that  chlorides  or  bromides  are  not  present  to 

TABLE  IV. 
ANALYSIS  OF  PURE  POTASSIUM  IODIDE. 


Exp. 

KI  taken. 

KI  found. 

Error. 

grm. 

gnn. 

grm. 

(1) 

0.0814 

0.0816 

0.0002+ 

(2) 

0.0814 

0.0813 

0.0001- 

(3) 

0.0814 

0.0805 

0.0009+ 

(4) 

0.0815 

0.0809 

0.0006- 

(5) 

0.0814 

0.0808 

0.0006- 

(6) 

0.0814 

0.0806 

0.0008- 

(7) 

0.0814 

0.0812 

0.0002- 

(8) 

0.1628 

0.1624 

0.0004- 

(9) 

0.1628 

0.1617 

0.0011- 

(10) 

0.1628 

0.1621 

0.0007- 

(11) 

0.1628 

0.1619 

0.0009- 

(12) 

0.1628 

0.1624 

0.0004- 

(13) 

0.1628 

0.1621 

0.0007- 

(14) 

0.1628 

0.1626 

0.0002- 

(15) 

0.2442 

0.2451 

0.0009+ 

(16) 

0.2442 

0.2442 

0.0000 

(17) 

0.2442 

0.2439 

0.0003- 

(18) 

0.3256 

0.3258 

0.0002+ 

(19) 

0.3256 

0.3256 

0.0000 

(20) 

0.3256 

0.3258 

0.0002+ 

(21) 

0.3256 

0.3272 

0.0016+ 

(22) 

0.3256 

0.3256 

0.0000 

(23) 

0.4071 

0.4076 

0.0005+ 

(24) 

0.4071 

0.4080 

0.0009+ 

(25) 

0.4071 

0.4073 

0.0002+ 

TO   THE  ANALYSIS  OF  IODIDES.  41 

any  considerable  extent.  For  determining  the  standard  of  a 
solution  of  nearly  pure  potassium  iodide,  employed  in  so  many 
laboratory  processes,  it  should  find  useful  application. 

In  Table  IV  are  comprised  a  number  of  experiments  made 
exactly  like  those  which  seemed  to  give  the  best  results  in 
the  series  of  Table  II.  The  iodide  and  an  excess  of  iodate 
(5  cm3  of  the  0.5  per  cent  solution  to  every  20  cm3  of  ^  iodide) 
were  made  to  interact  in  a  volume  of  about  150  cm3,  5  cm3  of 
sulphuric  acid  (1  :  3)  were  used  to  bring  about  the  reaction, 
10  cm3  of  potassium  bicarbonate  were  added  after  the  neutrali- 
zation of  the  sulphuric  acid  was  complete,  and  the  free 
iodine  was  estimated  by  titrating  decinormal  arsenious  acid, 
the  manipulation  being  like  that  previously  described  in 
detail. 

The  average  result  of  a  series  of  several  determinations  in 
which  a  great  excess  (0.1  grm.)  of  potassium  iodate  was 
used,  proved  to  be  practically  identical  with  that  of  a  similar 
series  in  which  only  a  small  excess  of  the  iodate  was  employed, 
so  that  it  appears  to  be  unnecessary  in  any  practical  work  to 
restrict  the  amount  of  iodate  below  the  amount  necessary  to 
decompose  the  maximum  quantity  of  potassium  iodide  which 
we  have  handled,  namely,  0.4  grm. 

It  appears  that  for  the  estimation  of  iodine  in  a  soluble 
iodide  free  from  notable  amounts  of  chlorides  or  bromides, 
this  method,  depending  as  it  does  upon  a  single  standard 
solution,  is  simple,  fairly  accurate,  and  rapid. 


vn 


THE  ACTION  OF   UREA  AND  PRIMARY  AMINES 
ON   MALEIC   ANHYDRIDE.* 

BY  FREDERICK  L.  DUNLAP  AND  ISAAC  K.  PHELPS. 

IN  a  former  article,  f  a  method  was  described  by  one  of  us 
for  obtaining  imides  by  the  action  of  urea  on  the  anhydrides 
of  dibasic  acids.  It  was  shown  that  the  formation  of  imides 
by  the  interaction  of  urea  and  anhydrides  was  to  be  explained 
by  the  addition  of  the  urea  and  the  anhydrides  to  form  an 
acid,  which  when  heated,  decomposed  with  the  formation  of 
the  imide,  carbon  dioxide  and  ammonia.  These  reactions  can 
be  shown  by  the  following  equations  : 


COOH 

.  ^O.^T 

<COOH  :     <CO> 

In  some  cases  this  intermediate  product,  formed  by  the 
addition  of  the  reacting  substances,  was  readily  isolated,  and 
upon  heating  to  certain  temperatures,  it  was  found  that  this 
addition-product  decomposed  with  the  formation  of  the  imide. 

It  was  found,  in  particular,  that  maleic  anhydride  formed 
an  addition-product  with  urea,  and  we  hoped  to  obtain  the 
unknown  imide  of  maleic  acid  by  the  decomposition  of  this 
addition-product. 

Male-uric  Acid.  —  When  equal  molecules  of  maleic  anhydride 
and  urea  are  heated  to  100°  -105°  C.  in  an  oil-bath,  the  mixture 
melts,  and  if  this  temperature  is  maintained  for  a  short  time, 

*  From  Am.  Chem.  Jour.,  xix,  492. 

t  Am.  Chem.  Jour.,  xviii,  332.    Volume  I,  355. 


ACTION  OF  UREA  AND  PRIMARY  AMINES.          43 

the  liquid  solidifies.  This  product  was  purified  by  crystalli- 
zation from  alcohol,  and  when  pure  was  perfectly  white  and 
melted  at  167.5°-168°  with  decomposition.  Upon  analysis 
the  following  results  were  obtained: 

0.2148   gram  substance   gave  0.2968   gram   C02  and  0.0751 
gram  H20. 

Calculated  for 


C  37.97  37.69 

H  3.80  3.88 

Maleiiric  acid  is  soluble  in  hot  water  and  alcohol,  fairly 
insoluble  in  cold  and  insoluble  in  chloroform,  benzene,  ligrom, 
carbon  disulphide,  acetone,  and  ether. 

Its  formation  and  structure  are  shown  in  the  following 
equation  : 

H-C-COV  H-C-CONHCONHa 

II  ^0  +  NH3CONH2=          || 

H-C-CCT  H-C-COOH 

No  difficulty  was  expected  in  the  formation  of  the  imide  of 
maleic  acid,  either  from  the  maleiiric  acid,  or  by  the  direct 
heating  of  maleic  anhydride  and  urea;  for,  in  the  cases 
already  studied,*  the  reaction  ran  with  great  smoothness  and 
ease.  But,  for  some  unknown  reason,  the  imide  of  maleic 
acid  could  not  be  obtained  —  at  least  in  quantities  for  complete 
identification.  A  great  many  experiments  were  carried  out 
under  varying  conditions,  but  only  under  one  set  of  conditions 
could  any  product  be  isolated,  and  then  unfortunately  the 
yield  was  so  distressingly  small  that  the  study  of  the  compound 
had  to  be  abandoned.  When  equal  molecules  of  maleic 
anhydride  and  urea  are  heated  in  a  boiling-flask,  at  first  slowly, 
and  finally  at  the  full  heat  of  a  Bunsen  burner,  a  vigorous 
evolution  of  carbon  dioxide  and  ammonia  took  place,  and  a 
small  amount  of  a  dark-colored  distillate  was  obtained,  which 
on  cooling  solidified  and  gave  a  melting-point  of  130.5°  after 
crystallization  from  anhydrous  acetone.  This  same  crystalline 

*  Loc.  cit. 


44  THE  ACTION  OF  UREA   AND  PRIMARY 

product  was  also  obtained  on  distilling  the  maleiiric  acid.  In 
both  cases  a  large  amount  of  carbonaceous  residue  remained 
in  the  boiling-flask. 

Although  the  method  of  formation  of  this  compound  is 
exactly  parallel  to  that  employed  in  the  preparation  of 
succinimide,*  yet  the  identity  of  this  product  cannot  be 
regarded  as  established  until  sufficient  quantities  of  it  have 
been  prepared  to  subject  it  to  elementary  analysis. 

The  Action  of  Primary  Amines  on  Maleic  Anhydride. 

Anschiitzf  in  1889  published  a  method  by  which  anilic  acids 
could  be  very  easily  prepared  by  dissolving  molecular  quantities 
of  aniline  and  the  anhydride  of  dibasic  acids  in  dry  chloroform, 
when  after  standing  for  a  short  tune  the  anilic  acid  crystallized 
out.  Following  out  this  line  and  method  of  formation, 
suggested  by  Anschiitz's  work,  some  derivatives  of  maleic 
acid  have  been  prepared  and  studied. 

p-Tolylmaleamic  Acid. — Molecular  proportions  of  maleic 
anhydride  and  ^>-toluidine  were  dissolved  in  dry  chloroform, 
and  upon  mixing  these  two  solutions,  there  immediately 
separated  out  a  light  yellow  precipitate.  After  standing  for  a 
time  to  insure  the  complete  precipitation  of  the  product,  it 
was  filtered  off  and  purified  by  crystallization  from  alcohol, 
from  which  it  separated  in  lemon-yellow  needles.  The  pure 
product  melted  at  201°,  with  evolution  of  gas.  It  is  readily 
soluble  in  ether,  acetone,  and  hot  alcohol,  but  insoluble  in 
benzene,  ligrom,  carbon  disulphide,  chloroform  and  water. 
Upon  analysis  the  f ollowing  results  were  obtained  : 

0.2072  gram  substance  gave  0.4881  gram  C02,  and  0.1068 
gram  H20. 

Calculated  for  v       ^ 

CUHU0SN.  Found' 

C  64.39  64.24 

H  5.37  5.73 

*  Loc.  cit.  t  Ber.  Dtsch.  chem.  Ges.,  xx,  3214 


AMINES   ON  MALEIC  ANHYDRIDE.  45 

The  structure  and  mode  of  formation  of  this  ^>-tolylmaleamic 
acid  may  be  seen  from  the  following  equation  : 


H-C-COV  XCH8         H~C-CONHC6H4CH8(p) 


C6H4 


N 


H-C-C(r  NH2(p)H-C-COOH 

o-Totylmaledmic  Acid.  —  This  acid  was  prepared  similarly  to 
the  ^-tolylmaleamic  acid,  using  molecular  proportions  of 
maleic  anhydride  and  0-toluidine.  The  product  was  purified 
by  crystallization  from  alcohol,  from  which  it  separated  in 
bunches  of  thick  light  yellow  prisms.  When  pure  it  melted 
at  117.5°-118°.  It  is  readily  soluble  hi  acetone,  and  in 
alcohol;  sparingly  soluble  in  chloroform,  and  insoluble  in 
ligroin,  carbon  disulphide,  benzene,  water,  and  ether.  Analy- 
sis gave  the  following  results  : 

0.2124  gram  substance  gave  0.4993  gram  C02,  and  0.1146 
gram  H20. 


C  64.39  64.12 

H  5.37  5.99 

It  has  the  structure  represented  by  the  formula 

H  -  C  -  CONHC6H4CH8  (o) 

II 
H-C-COOH 

p-Naphthylmaleamic  Acid.  —  When  molecular  quantities  of 
^-naphthylamine  and  maleic  anhydride  were  mixed  in  dry 
chloroform  solution  they  reacted  and  united  as  readily  as  did 
the  o-  and  ^>-toluidines  and  the  anhydride.  Almost  immediately 
a  yellow  crystalline  precipitate  separated  out.  This,  after 
standing  a  short  time,  was  filtered  off  and  crystallized  from 
alcohol.  It  separated  from  alcohol  in  small  bright  yellow 
needles,  which,  when  pure,  melted  at  200°  with  evolution  of 
gas.  It  is  soluble  in  acetone  and  alcohol,  but  insoluble  in 
ether,  carbon  disulphide,  chloroform,  benzene,  ligroin,  and 
water.  Analysis  gave  the  following  results  : 


46         ACTION  OF  UREA   AND  PRIMARY  AMINES. 

0.2033   gram  substance  gave  0.5226   gram  C02  and  0.0948 
gram  H20. 

Calculated  for  v       , 

C14HU03N.  FoTmd' 

C  69.71  70.11 

H  4.56  5.18 

This  addition-product  is  obviously  formed  as  follows : 
H  _  C  -  COV  H-C-  CONHC10H7 


0  +  C10H7NH2  = 


H-C-CCT  H-C-COOH 

a-Naphthylamine  appears  also  to  be  added  readily  to 
maleic  anhydride,  but  the  product  has  not  been  submitted 
to  analysis. 


VIII 


THE  SEPARATION  OF  ALUMINUM  AND  BERYL- 
LIUM BY  THE  ACTION  OF  HYDROCHLORIC 
ACID. 

BY  FRANKE  S.  HAVENS* 

IN  a  former  paper  f  a  method  was  described  for  the  determi- 
nation of  aluminum  in  the  presence  of  iron,  based  upon 
the  fact  that  the  hydrous  aluminum  chloride  A1C13 .  6H2O 
is  practically  insoluble  in  a  mixture  of  strong  hydrochloric 
acid  and  anhydrous  ether  saturated  with  hydrochloric  acid 
gas,  while  the  ferric  chloride  is  entirely  soluble  in  that 
medium. 

The  work  to  be  described  in  this  paper  is  an  extension 
of  this  process  to  cover  the  separation  of  aluminum  from 
beryllium,  with  the  subsequent  determination  of  the  beryllium 
by  weighing  as  the  oxide  after  conversion  to  the  nitrate  and 
ignition. 

The  aluminum  chloride  solution  was  prepared  by  dissolving 
the  so-called  pure  aluminum  chloride  of  commerce  in  as  little 
water  as  possible,  precipitating  and  washing  free  from  iron 
with  strong  hydrochloric  acid,  dissolving  the  chloride  thus 
obtained  in  water,  precipitating  the  hydroxide  by  ammonia, 
washing  the  precipitate  free  from  all  alkalies,  and  redissolving 
it  in  hot  hydrochloric  acid.  From  this  solution,  after  cool- 
ing, gaseous  hydrochloric  acid  precipitated  the  pure  hydrous 
chloride.  This  prepared  chloride  was  dissolved  in  water 
and  the  solution  standardized  by  precipitating  with  ammonia 
the  hydroxide  from  weighed  portions  and  weighing  as  the 
oxide.  The  solution  of  beryllium  used  was  made  by  dis- 

*  From  Am.  Jour.  Sci.,  iv,  111. 

t  Gooch  and  Havens,  Am.  Jour.  Sci.,  ii,  416.    This  volume,  p.  20. 


48       SEPARATION  OF  ALUMINUM  AND  BERYLLIUM 


solving  in  water  beryllium  chloride  found  to  be  free  from 
iron  by  the  sulphocyanate  test,  and  giving  no  precipitate 
when  tested  by  the  gaseous  hydrochloric  acid  process  to  be 
described  later  on.  This  was  standardized  by  precipitating 
with  ammonia  the  hydroxide  from  weighed  portions  and 
weighing  the  ignited  oxide  in  the  usual  manner. 

In  the  experiments  of  Table  I,  weighed  portions  of  the 
aluminum  solution  were  mixed  with  portions  of  the  beryllium 
chloride  solution  representing  from  0.01  gram  to  0.10  gram 
of  the  oxide,  an  equal  volume  of  a  mixture  of  strong  hydro- 
chloric acid  and  ether  (taken  in  equal  parts)  was  added  to  the 
solution  of  the  mixed  chlorides,  and  the  whole  was  completely 
saturated  with  gaseous  hydrochloric  acid  while  kept  at  a 
temperature  of  about  15°  C.  by  immersing  the  receptacle 

TABLE  I. 


Bxp. 

A12O8  taken  in 
solution  as  the 
chloride. 

found. 

Final 
volume. 

Error. 

(1) 

grm. 

0.1046 

gnn. 
0.1044 

cms 
12 

grm. 
0.0002- 

(2) 

0.1046 

0.1038 

12 

0.0008- 

(3) 

0.1067 

0.1066 

12 

0.0001- 

(4) 

0.1071 

0.1063 

12 

0.0008- 

(5) 

0.1059 

0.1064 

30 

0.0005- 

in  running  water.  Ether  was  added,  equal  in  volume  to 
the  aqueous  aluminum  and  beryllium  solutions  originally 
taken,  and  the  current  of  gas  again  turned  on  until  satura- 
tion was  complete.  By  this  treatment  there  is  present  at  the 
end  of  the  saturation  a  volume  of  ether  equal  to  that  of 
the  aqueous  hydrochloric  acid  introduced  and  generated. 
The  finely  crystalline  precipitate  of  aluminum  chloride  was 
caught  on  asbestos  in  a  filter  crucible  washed  with  a  pre- 
viously prepared  mixture  of  hydrochloric  acid  and  ether  in 
equal  parts  saturated  at  15°  C.  with  hydrochloric  acid  gas, 
and  dried  for  half  an  hour  at  a  temperature  of  150°  C.  It 
was  next  covered  with  a  layer  of  pure  mercuric  oxide,  which 
had  been  tested  and  found  to  leave  no  residue  on  volatilizing, 


BY  THE  ACTION  OF  HYDROCHLORIC  ACID. 


49 


and  the  crucible  was  gently  heated  over  a  low  flame  under 
a  ventilating  hood  and  finally  ignited  over  the  blast. 

From  these  results  it  is  obvious  that  the  aluminum  chloride 
may  be  determined  in  the  presence  of  beryllium  chloride  with 
reasonable  accuracy. 

The  beryllium  may  be  recovered  in  the  filtrate  from  the 
aluminum  chloride  by  precipitation  with  ammonia  after  nearly 
complete  evaporation  of  the  acid.  It  was  found,  however, 
upon  trial  that  the  conversion  of  the  chloride  to  the  oxide 
without  precipitation  and  filtration  may  be  easily  accomplished 
by  treatment  with  nitric  acid  and  ignition.  The  results  of 
Table  II  indicate  this  clearly.  In  these  experiments  weighed 

TABLE  II. 


BeO  taken  in 

Exp. 

solution  as  the 

BeO  found. 

Error. 

chloride. 

grin. 

gnu. 

grm. 

0) 

0.0483 

0.0481 

0.0002- 

2) 

0.0483 

0.0483 

0.0000 

(3) 

0.1076 

0.1085 

0.0009+ 

portions  of  the  beryllium  solution  were  evaporated  just  to 
dryness  on  a  radiator,  care  being  taken  not  to  heat  to  the 
volatilizing  point  of  the  beryllium  chloride,  a  few  drops  of 
strong  nitric  acid  were  added,  the  liquid  was  evaporated, 
and  the  residue  heated  —  at  first  gently,  to  break  up  the 
nitrate  safely  and  finally  on  the  blast.  It  was  found  that 
this  conversion  of  the  beryllium  to  the  nitrate  can  be  carried 
on  in  platinum  without  attacking  that  metal  appreciably, 
providing  care  be  taken  to  remove  thoroughly  all  excess  of 
hydrochloric  acid  before  the  nitric  acid  is  added  to  the  dry 
residue. 

In  Table  III,  (l)-(9),  are  given  the  results  of  experiments 
in  which  both  the  aluminum  and  the  beryllium  were  deter- 
mined —  the  former  by  precipitation  as  the  hydrous  chloride 
and  weighing  as  the  oxide  after  igniting  with  mercuric  oxide : 
the  latter  by  the  conversion  of  the  chloride,  through  the 

VOL.    II.  —  4 


50       SEPARATION  OF  ALUMINUM  AND  BERYLLIUM 


nitrate,  into  the  oxide.  In  experiment  (10)  (made  to  get  a 
comparison  of  the  methods)  the  beryllium  was  recovered 
by  precipitating  the  hydroxide  with  ammonia  from  the  par- 
tially evaporated  solution  of  the  chloride  after  removing  the 
aluminum. 

In  experiments  (1)  to  (5),  inclusive,  the  aluminum  was 
determined  exactly  as  previously  described ;  in  (6)  and  (7) 
the  solutions  (being  originally  larger)  were  concentrated 
by  evaporation  previous  to  the  addition  of  the  ether  and 
hydrochloric  acid  mixture.  In  experiments  (8),  (9)  and 
(10),  the  treatment  was  varied  advantageously  by  saturating 
the  aqueous  solution  directly  with  hydrochloric  acid  gas 
before  adding  an  equal  volume  of  ether,  and  completing  the 
saturation. 

TABLE  III. 


A1203  taken 

BeO  taken 

Exp. 

in  solution 
as  the 

AM>, 

Error. 

Final 
volume. 

in  solution 
as  the 

BeO 

found. 

Error. 

chloride. 

chloride. 

grin. 

grm. 

grm* 

cm3 

gnu* 

grin. 

grm. 

(1) 

0.1059 

0.1058 

0.0001- 

12 

0.0198 

0.0204 

0.0006+ 

(2) 

0.1053 

0.1044 

0.0009- 

12 

0.0194 

0.0196 

0.0002+ 

(3) 

0.1065 

0.1059 

0.0006- 

12 

0.0197 

0.0205 

0.0008+ 

(4 

0.1068 

0.1060 

0.0008- 

12 

0.0199 

0.0207 

0.0008+ 

(5) 

0.1049 

0.1047 

0.0002- 

12 

0.0198 

0.0208 

0.0010+ 

(6) 

0.1060 

0.1057 

0.0003- 

12 

0.0977 

0.0969 

0.0008- 

(7) 

0.1064 

0.1063 

0.0001- 

12 

0.1085 

0.1084 

0.0001- 

(8) 

0.1046 

0.1038 

0.0008- 

30 

0.1083 

0.1087 

0.0004+ 

(9) 

0.1051 

0.1048 

0.0003- 

30 

0.1071 

0.1078 

0.0007+ 

(10) 

0.1076 

0.1075 

0.0001- 

30 

0.1086 

0.1094 

0.0008+ 

These  results  are  plainly  very  good. 

The  manipulation  of  the  process  is  not  difficult.  The 
gaseous  hydrochloric  acid  is  most  conveniently  produced  by 
the  well  known  method  of  treating  with  strong  sulphuric 
acid  in  regulated  current  a  mixture  of  strong  aqueous 
hydrochloric  acid  and  common  salt.  A  platinum  dish  hung 
in  an  inverted  bell-jar,  provided  with  inlet  and  outlet  tubes 
through  which  the  current  of  water  for  cooling  is  passed, 
makes  the  best  container  for  the  solution  to  be  saturated 
with  the  gas.  It  is  advantageous  to  arrange  the  filtration 


BY  THE  ACTION  OF  HYDROCHLORIC  ACID.          51 

upon  asbestos  so  that  the  filtrate  and  washings  may  be 
caught  directly  in  the  crucible  (placed  under  the  bell-jar  of 
the  filter  pump)  in  which  the  subsequent  evaporation  is  to 
be  effected.  The  heating  of  the  strong  acid  solution  must 
be  gradual  and  conducted  with  care  to  prevent  mechanical 
loss  by  a  too  violent  evolution  of  the  gaseous  acid. 


IX 

THE  TITRATION   OF  SODIUM  THIOSULPHATE 
BY  IODIC  ACID. 

BY  CLAUDE  F.  WALKER .* 

THIS  investigation  was  undertaken  to  determine  the  nature 
and  limitations  of  the  reaction  between  iodic  acid  and  thiosul- 
phuric  acid,  and  to  show  the  expediency  of  employing  iodic 
acid  in  standard  solution  for  the  direct  titration  of  sodium 
thiosulphate.  Rieglert  states  that  iodic  acid  is  readily  ob- 
tained in  the  pure  state,  that  it  may  be  accurately  weighed 
out,  and  that  a  solution  of  it  may  be  exactly  made  up  to  a 
desired  strength  and  kept  for  a  long  tune  unaltered.  He 
further  states  that  when  a  solution  of  sodium  thiosulphate  is 
titrated  with  iodic  acid  the  reaction  takes  place  according  to 
the  equation, 

6Na2S203  +  6HI03  =  3Na2S<06  +  5NaI03  +  Nal  +  3H20, 

under  which  circumstances  no  free  iodine  will  be  evolved 
until  all  the  sodium  thiosulphate  has  been  oxidized  to  tetra- 
thionate ;  the  first  drop  of  iodic  acid  in  excess,  however,  will 
react  with  the  sodium  iodide  that  has  been  formed,  and  sepa- 
rate iodine,  as  shown  by  the  equation, 

5NaI  +  6HI03  =  5NaI08  +  3H20  +  3I2, 

thus  furnishing  an  accurate  means  for  determining  the  end 
point. 

A  careful  repetition  of  the  work  of  Riegler  has  shown  that 
his  conclusions  are  in  a  large  measure  erroneous.  Thus,  it 
has  been  found  that  the  ordinary  "  chemically  pure  "  iodic 
acid,  purchased  from  reliable  manufacturers,  is  likely  to  con- 

*  From  the  Amer.  Jour.  Sci.,  iv,  235. 

t  Riegler,  Zeitschr.  anal.  Chem.,  xxxr,  308. 


TITRATION  OF  SODIUM  THIOSULPHATE,  ETC.         53 

tain  more  than  the  theoretical  amount  of  iodine,  due  probably 
to  the  presence  of  the  anhydride,  although  iodic  acid  can  be 
safely  employed  for  standardizing  when  it  is  made  in  the 
laboratory  by  dissolving  the  purified  anhydride,  crystallizing 
out  the  acid,  and  drying  over  sulphuric  acid.  Such  a  care- 
fully prepared  product,  if  used  immediately,  will  be  found  to 
contain  the  theoretical  amount  of  iodine.  Riegler's  proposed 
method  of  titration  depends  on  two  different  reactions,  and  to 
insure  the  accuracy  of  the  process  these  must  be  definite,  com- 
plete and  non-reversible  under  the  conditions  of  analysis. 
Thus  one  molecule  out  of  every  six  of  iodic  acid  should  be 
reduced  by  six  molecules  of  thiosulphate,  with  the  formation 
of  a  neutral  mixture  of  iodide  and  iodate,  free  from  other 
oxidizing  or  reducing  substances.  Under  these  circumstances 
it  might  be  expected  that  iodine  will  be  liberated  by  the  first 
trace  of  iodic  acid  hi  excess.  It  has  been  found  by  investiga- 
tion, however,  that  although  the  main  reaction  between  iodic 
acid  and  sodium  thiosulphate  may  result  in  the  formation  of 
sodium  tetrathionate  in  the  proportions  given,  there  is  never- 
theless striking  evidence  of  some  other  obscure  action  of  the 
thiosulphate,  which  influences  the  reduction  of  the  iodic  acid 
in  such  a  way  as  to  make  it  impossible  to  calculate  the  analy- 
ses according  to  Riegler's  reaction.  Moreover,  a  peculiar 
"  after-coloration  "  which  invariably  follows  the  first  formation 
of  the  starch  blue  during  the  titration  of  one  solution  against 
the  other,  seems  to  point  to  the  possibility  that  the  reaction 
between  the  iodide  and  iodic  acid  is  dependent,  under  these 
circumstances,  on  conditions  of  time  and  mass  for  its  com- 
pleteness. It  is  not  impossible  that  some  third  compound  of 
iodine  unstable  in  its  nature,  may  be  formed  as  an  interme- 
diate product  and  thus  delay  the  liberation  of  iodine.  In  con- 
sideration of  the  results  that  have  been  obtained  it  appears 
that  Riegler's  proposed  process  for  standardizing  sodium  thio- 
sulphate, as  well  as  his  related  method  for  the  analysis  of 
iodides,*  must  remain  impracticable  unless  they  can  be  modi- 
fied so  as  to  do  away  with  a  number  of  sources  of  error. 

*  Riegler,  Zeitschr.  anal.  Chem.,  xxxv,  305. 


54  TITRATION  OF  SODIUM  THIOSULPHATE 

The  analyses  of  solutions  of  iodic  acid,  during  the  entire 
course  of  the  work,  was  invariably  performed  by  adding  to 
the  portion  of  the  solution  to  be  analyzed  an  excess  of  potas- 
sium iodide,  acidifying  with  5  cm3  of  dilute  (1  :  3)  sulphuric 
acid,  and  recovering  the  liberated  iodine  by  directly  titrating 
the  acid  solution  with  sodium  thiosulphate,  or  by  neutralizing 
with  potassium  bicarbonate  in  excess,  and  directly  titrating 
the  alkaline  solution  with  arsenious  acid.  In  the  latter  case 
the  neutralization  was  performed  in  a  trapped  Drexel  wash- 
ing bottle  such  as  has  been  described  in  connection  with  the 
analysis  of  iodides.*  In  either  case  one-sixth  of  the  iodine 
recovered  was  calculated  to  iodic  acid,  according  to  the  terms 
of  the  equation, 

SHI  +  HI03  =  31,  +  3H20. 

It  follows  from  these  proportions  that  to  bring  the  analyses 
within  the  range  of  the  decinormal  solutions  ordinarily 
employed,  the  iodic  acid  taken  for  analysis  must  be  restricted 
to  comparatively  small  amounts.  In  the  present  work  it  was 
found  convenient  to  analyze  the  iodic  acid  in  quantities  not 
much  exceeding  one-tenth  of  a  gram,  in  which  case  the  varia- 
tion in  the  results  in  the  same  series  is  found  to  be  almost  inap- 
preciable. In  both  variations  of  the  process  one  or  two  blank 
analyses  were  invariably  made,  by  performing  the  operation  as 
detailed,  except  that  no  iodic  acid  was  employed,  and  the  cor- 
rection of  one  drop  of  iodine  thereby  shown  to  be  necessary  to 
bring  out  the  starch  blue  was  uniformly  applied  in  the  ana- 
lytical work. 

To  determine  whether  or  not  the  purity  of  the  ordinary 
iodic  acid  is  sufficient  to  admit  of  its  direct  application  in 
standard  solutions,  a  series  of  experiments  was  made.  Two 
different  samples  of  "  chemically  pure  "  iodic  acid  were  used. 
The  first  was  in  coarse  granular  crystals,  and  the  second  was 
in  the  form  of  fine  powder.  Quantities  of  both  of  these 
were  dried  in  a  desiccator  over  sulphuric  acid  to  constant 
weight.  Neither  sample  lost  weight  appreciably  when  left 

*  Gooch  and  Walker,  Am.  Jour.  Sci.  iii,  293.    This  volume,  p.  33. 


BY  IODIC  ACID. 


55 


for  a  considerable  time  on  the  scale  pan.  A  third  sample  of 
iodic  acid  was  prepared  by  dissolving  a  quantity  of  the 
purest  obtainable  iodic  anhydride  in  water,  and  evaporating  at 
ordinary  temperature.  The  resulting  crystalline  mass  was 
dried  over  sulphuric  acid  in  a  desiccator  for  one  week,  until  it 
ceased  to  lose  weight,  when  it  was  presumed  to  consist  of  the 
pure  normal  acid.  Two  presumably  decinormal  solutions  of 
each  of  the  first  two  samples,  and  one  such  solution  of  the 
third  sample  of  iodic  acid  were  made  by  weighing  out  17.585 
grms.  and  dissolving  in  exactly  one  liter  of  water  at  15°  C. 
Convenient  portions  of  each  of  these  solutions  were  analyzed 
in  the  manner  described,  with  results  shown  in  the  following 
table,  averaged  from  many  determinations. 

TABLE  I. 
ANALYSES  OF  APPROXIMATELY  ^  IODIC  ACID. 


Solution 
analyzed. 

Sample 
used. 

mO3  taken. 

HIOS  found. 

Error. 

grin. 

grm. 

grm. 

I 

A 

0.1055 

0.1066 

0.0011+ 

II 

A 

0.1055 

0.1062 

0.0007+ 

III 

B 

0.1055 

0.1065 

0.0010+ 

IV 

B 

0.1055 

0.1073 

0.0018+ 

V 

C 

0.1055 

0.1053 

0.0002- 

These  results  show  that  while  the  deviation  from  the 
theoretical  strength  of  the  solution  in  the  case  of  the  acid 
prepared  from  the  anhydride  is  hardly  appreciable,  and  will 
not  affect  the  accuracy  of  any  work  in  which  the  solution 
may  be  applied  as  a  means  of  standardization,  the  solutions 
made  from  the  purchased  product,  on  the  other  hand,  contain 
a  very  appreciable  amount  of  iodine  in  excess  of  the  theoretical. 
That  iodic  acid  is  somewhat  unstable  at  30-40°  C.,  gradually 
losing  water  with  the  formation  of  the  anhydride,*  is  well 
known,  and  it  is  quite  possible  that  to  some  such  gradual 
change  as  this  must  be  attributed  the  fact  that  the  ordinary 
iodic  acid  cannot  be  safely  employed  as  a  means  of  standard- 
ization unless  its  purity  be  directly  determined  by  analysis. 

*  Dammer,  Anorg.  Chem.,  i,  564. 


56 


TITRATION  OF  SODIUM  THIOSULPHATE 


To  determine  whether  a  solution  of  iodic  acid,  once  prepared 
and  standardized,  will  retain  its  strength  for  a  long  period  of 
time,  two  such  solutions  were  kept  for  four  months  (in  the 
dark)  and  then  again  analyzed.  The  results  (averages  of 
several  determinations),  given  in  Table  II,  substantiate  the 
observation  of  Riegler  that  a  solution  of  iodic  acid  will  remain 
of  constant  strength. 

TABLE  II. 
CONSTANCY  OF  STRENGTH  OF  IODIC  ACID  SOLUTIONS. 


Iodic  acid 
solution. 

First  analysis. 
HI03  found. 

Second  analysis 
(after  four  months). 
HIO3  found. 

Variation. 

I 
II 

grm. 
0.1073 
0.1049 

grm. 

0.1072 
0.1046 

grm. 
0.0001- 
0.0003- 

An  approximately  one-twentieth  normal  solution  of  "  chemi- 
cally pure"  sodium  thiosulphate  was  made  and  its  exact 
strength  ascertained  by  titrating  it  with  standardized  iodine. 
A  series  of  analyses  made  by  oxidizing  the  sodium  thiosulphate 
to  sulphate,  and  precipitating  and  weighing  as  barium  sulphate, 
gave  results  identical  with  those  obtained  with  iodine,  proving 
that  all  the  sulphur  present  in  the  solution  was  in  the  form 
of  thiosulphate.  According  to  Riegler's  equation,  sodium 
thiosulphate  and  iodic  acid  react  molecule  for  molecule,  and 
solutions  of  these  substances  should  therefore  require  for 
their  mutual  saturation  volumes  inversely  proportional  to 
their  concentration.  It  was  found,  however,  that  when  the 
one-twentieth  normal  solution  of  sodium  thiosulphate  that  has 
been  described  was  titrated  in  the  presence  of  starch  emulsion 
with  an  approximately  decinormal  solution  of  iodic  acid, 
prepared  from  the  anhydride,  a  distinctly  blue  color  was 
produced  long  before  the  theoretical  amount  of  iodic  acid  had 
been  added.  It  was  further  noticed  that  the  end-point  of  the 
reaction  was  far  from  distinct,  a  faint  tinge  of  blue  at  first 
being  visible,  then  suddenly  becoming  deeper,  and  immediately 
reappearing  when  bleached  with  sodium  thiosulphate.  The 


BY  IODIC  ACID.  57 

deficiency  in  the  amount  of  iodic  acid  actually  required  to 
produce  the  blue  color  was  not  lessened  by  introducing  only 
three-fourths  of  the  theoretical  amount  of  iodic  acid,  and 
estimating  the  residual  thiosulphate  with  iodine.  It  was 
found,  however,  that  the  addition  of  a  considerable  quantity 
of  potassium  iodide  to  the  solution,  either  before  or  during 
the  titration,  had  the  marked  effect  of  making  the  reaction 
sharp  and  distinct,  entirely  preventing  the  "  after  separation  " 
of  iodine,  at  the  same  time  postponing  the  appearance  of  the 
starch  blue  until  a  quantity  of  iodic  acid  had  been  added 
considerably  in  excess  of  the  theoretical.  These  experiments 
were  repeated  with  entirely  different  reagents,  and  under 
varied  conditions  of  concentration,  the  results  in  every  case 
exactly  confirming  those  already  observed. 

For  the  purpose  of  more  particularly  investigating  this 
subject,  there  were  prepared  and  standardized  an  approxi- 
mately decinormal  solution  of  sodium  thiosulphate,  and  an 
approximately  one-fiftieth  normal  solution  of  iodic  acid. 
Measured  portions  of  the  sodium  thiosulphate  solution  were 
titrated  with  the  iodic  acid  in  the  presence  of  starch  emulsion 
under  varying  conditions  of  mass,  time  and  dilution. 

To  determine  the  variability  of  the  end-point  of  the  reaction 
when  the  titration  was  conducted  as  directed  by  Riegler,  a 
series  of  experiments  was  made.  Measured  amounts  of  the 
sodium  thiosulphate  'solution  were  drawn  from  a  burette  into 
an  Erlenmeyer  beaker  of  suitable  capacity,  the  sides  of  the 
beaker  were  carefully  washed  down  with  a  small  amount  of 
water,  5  cm3  of  starch  emulsion  were  added,  and  the  iodic  acid 
was  slowly  dropped  into  the  small  bulk  of  acid  and  starch, 
with  constant  agitation  of  the  mixture,  until  the  first  tint  of 
blue  coloration  appeared.  The  results  obtained  are  given  in 
Table  III. 

These  experiments  indicate  that  the  constancy  of  the  end 
reaction  in  different  titrations  of  equal  volumes  of  the  same 
solution  depends  to  a  certain  degree  on  the  volume  of  sodium 
thiosulphate  taken.  The  results  in  the  case  of  the  maximum 
amounts  vary  within  a  range  of  1.04  cm3,  which  corresponds 


58 


TITRATION  OF  SODIUM  THIOSULPHATE 


TABLE  III. 

VARIATION  OF  THE  END  REACTION  BETWEEN  ~  SODIUM  THIOSULPHATE 
AND       lODIC  ACID,  IN  THE  ABSENCE  OP  POTASSIUM  IODIDE. 


Eip. 

NaAO, 
taken. 

HIO, 
introduced. 

Mean 
value. 

Variation. 

1) 

cm8 
6 

cm8 
28.131 

cm8 

cm8 
0.19- 

2) 

6 

27.79 

0.53- 

3) 

6 

28.03 

0.29- 

1 

6 
6 

28.32 
28.32 

28.32 

0.00 
0.00 

6) 

6 

28.71 

0.39+ 

7) 

6 

28.83 

0.51+ 

8) 

6 

28.43 

0.11+ 

(9) 

4 

18.94  1 

0.26+ 

(10) 
(11) 

4 

4 

18.67  1 
18.60  f 

18.68 

0.01- 
0.18— 

(12) 

4 

18.60  J 

0.08- 

to  0.0035  grm.  of  iodic  acid,  while  the  average  variation  is 
0.25  cm3,  corresponding  to  0.0009  grm.  The  variation  in  the 
analyses  of  the  smaller  amounts  is  less,  the  range  being  0.44 
cm3,  corresponding  to  0.0015  grm.,  and  the  average  variation 
being  0.13  cm3,  or  0.0005  grm.  The  probable  error  which 
these  irregularities  would  introduce  in  any  series  of  practical 
analyses  by  this  method  is  obviously  greater  than  can  ordinarily 
be  permitted  in  iodometric  work. 

TABLE  IV. 
VARIATION  OP  THE  END  REACTION  BETWEEN  ^  SODIUM  THIOSULPHATE 

AND          lODIC    ACID,    IN   THE   PRESENCE    OF   POTASSIUM    IODIDE. 


Eip. 

Na2S203 
taken. 

HIOS 
introduced. 

Mean 
value. 

Variation. 

(1) 

cm8 
6 

cm8 
32.53  ] 

cm8 

cm8 
\  0.05+ 

(2) 

6 

32.45 

0.03- 

(3) 
(4) 

6 
6 

32.67 
32.37 

32.48 

• 

0.19+ 
0.11- 

(5) 

6 

32.36 

0.12- 

(6) 

6 

32.50 

0.02+ 

(7) 

4 

22.30  ' 

0.11+ 

(8) 
(9) 

4 
4 

21.98 
22.17 

22.19 

i 

0.21- 
0.02- 

(10) 

4 

22.30  J 

0.11+ 

BY  IODIC  ACID.  59 

The  experiments  detailed  in  Table  IV  were  performed 
exactly  similarly  to  those  of  the  last  series  except  that  two 
grams  of  potassium  iodide  were  added  to  the  sodium  thiosul- 
phate  before  the  titration  was  commenced. 

These  experiments  indicate  plainly  that  in  the  presence  of 
potassium  iodide  the  end  reaction  of  different  titrations  of 
equal  volumes  of  the  same  solution  is  practically  independent 
of  the  amount  taken  for  analysis.  The  results  in  the  case  of 
the  maximum  amounts  vary  within  a  range  of  0.31  cm3,  or 
0.0011  grm.  of  iodic  acid,  while  the  average  variation  is  0.09 
cm3,  corresponding  to  0.0003  grm.  The  variation  in  the 
analyses  of  the  smaller  amounts  is  practically  the  same  as  that 
of  the  larger,  the  range  being  0.32  cm3,  corresponding  to  0.0011 
grm.,  and  the  average  variation  being  0.11  cm3,  or  0.0004  grm. 
It  is  therefore  evident  that  the  presence  of  potassium  iodide  in 
the  sodium  thiosulphate  to  be  titrated  will  bring  the  variation 
of  the  formation  of  the  reading  tint  within  permissible  limits. 

A  series  of  experiments  was  made  to  determine  the  nature 
and  effect  of  the  "  after  coloration"  observed  to  take  place 
when  a  solution  of  sodium  thiosulphate,  free  from  potassium 
iodide,  was  titrated  with  iodic  acid  to  blue  coloration,  and  then 
bleached  with  sodium  thiosulphate.  The  titrations  were 
performed  in  the  usual  manner  except  that  the  volume  was 
adjusted  just  before  the  addition  of  the  iodic  acid,  and  the 
iodine  that  was  set  free  after  the  formation  of  the  first  reading 
tint  was  destroyed  at  fixed  intervals  with  measured  amounts  of 
sodium  thiosulphate.  The  results  are  given  in  Table  V. 

In  the  experiments  with  small  volumes  the  evolution  of 
iodine  in  any  considerable  quantity  ceased  after  two  or  three 
hours,  although  the  solution  would  become  recolored  as  often 
as  it  was  bleached  for  a  number  of  days.  The  traces  of  iodine 
thus  set  free,  however,  were  seldom  equivalent  to  more  than 
one  or  two  drops  of  sodium  thiosulphate.  The  larger  volumes, 
however,  continued  to  separate  iodine  in  abundance  for  a  very- 
long  time.  The  amount  of  iodine  thus  liberated  after  the  first 
coloration  evidently  varies  with  the  amount  of  iodic  acid 
required  for  the  titration,  although  not  strictly  proportional  to 


60 


TITRATION  OF  SODIUM  THIOSULPHATE 


TABLE  V. 
EFFECT  OF  DILUTION  AND  LAPSE  OF  TIME  ON  THE  "AFTER  COLORATION." 


„ 

^03 

HIO8 

intwi 

Na,S2Os  introduced. 

taken. 

intro- 
duced. 

15m. 

45m. 

lh.45m. 

2h.45in. 

20  h. 

Total. 

Volume. 

cm3 

cm8 

cm3 

cm3 

cm3 

cm3 

cm3 

cm8 

cm3 

(1) 

6 

27.68 

0.25 

0.13 

0.08 

0.00 

0.03 

0.49 

50 

(2) 

6 

27.70 

0.20 

0.10 

0.03 

0.03 

0.03 

0.39 

50 

'3) 

6 

28.17 

0.16 

0.10 

0.03 

0.01 

none. 

0.30 

50 

(4) 

6 

27.03 

0.60 

0.26 

0.09 

0.03 

none. 

0.98 

150 

(5) 

6 

27.60 

0.93 

0.28 

0.06 

0.04 

0.04 

1.35 

150 

6) 

6 

28.60 

1.34 

0.46 

0.17 

0.03 

0.14 

2.14 

200 

7) 

6 

28.85 

1.20 

0.50 

0.28 

0.06 

0.27 

231 

200 

8) 

6 

31.63 

1.46 

0.74 

0.10 

0.21 

0.23 

2.74 

250 

9) 

6 

29.90 

1.04 

0.60 

0.23 

0.15 

0.46 

2.48 

250 

( 

0) 

6 

36.09 

1.60 

1.23 

0.63 

0.34 

0.18 

3.98 

300 

11) 

6 

37.59 

1.65 

1.33 

0.72 

0.27 

0.10 

4.07 

300 

(12) 

6 

37.23 

1.92 

1.05 

0.64 

0.33 

* 

•  • 

300 

it.  Both  of  these  quantities  increase  at  a  regular  rate  with  the 
volume  of  the  solution. 

To  show  with  what  accuracy  the  reaction  between  sodium 
thiosulphate  and  iodic  acid  may  be  applied  to  the  direct 
estimation  of  one  of  these  substances  by  the  other,  the 
averaged  results  of  a  large  number  of  titrations  are  compared 
in  Table  VI.  The  operations  were  conducted  as  directed  by 
Riegler,  equal  measured  volumes  of  standardized  sodium 
thiosulphate  being  titrated  with  iodic  acid  of  known  strength, 
in  the  presence  of  starch  and  under  different  conditions  of 
time,  dilution,  and  mass,  the  volume  of  iodic  acid  required  to 
produced  the  blue  coloration  being  in  each  case  compared  with 
the  volume  theoretically  required  by  the  terms  of  Riegler's 
equation. 

These  results  show  plainly  that  the  amount  of  iodic  acid 
required  to  decompose  a  given  amount  of  sodium  thiosulphate 
may  be  considerably  above  or  below  that  required  by  the  terms 
of  Riegler's  equation.  Thus,  with  small  volumes,  and  in  the 
absence  of  potassium  iodide,  the  thiosulphate  is  destroyed  and 
the  separation  of  iodine  commences  when  only  93  per  cent  of 
the  theoretical  amount  of  acid  has  been  titrated.  At  higher 
*  No  observation. 


BY  IODIC  ACID. 


61 


TABLE  VI. 
TITRATION  OP       SODIUM  THIOSULPHATB  WITH 


IODIC  ACID. 


Exp. 

NaAO. 

taken. 

HIOS  used. 

HI08 

required 
by  theory. 

Error. 

Error. 

KI 

present. 

Volume. 

cm3 

cm3 

cm3 

cm8 

per  cent. 

grin. 

cm3 

(1) 

4 

f  18.68 

20.32 

1.64- 

8.0- 

50 

(2) 

6 

28.32 

30.48 

2.16- 

7.0- 

60 

(3) 

6 

27.32 

30.48 

3.16- 

10.0- 

150 

(4) 

6 

3fc  • 

28.73 

30.48 

1.75- 

6.0- 

200 

(5) 

6 

30.77 

30.48 

0.29+ 

0.01+ 

250 

(6) 

6 

36.97 

30.48 

6.49+ 

21.0+ 

300 

(7) 

6 

27.46 

30.48 

3.02- 

10.0- 

50 

8 

6 

26.15 

30.48 

4.33- 

14.0- 

150 

(9) 
(10) 

6 
6 

t" 

26.50 
28.16 

30.48 
30.48 

3.98- 
2.32- 

13.0- 
8.0- 

200 
250 

(11) 

6 

1  32.93 

30.48 

2.45+ 

8.0+ 

300 

(12) 

4 

.i  22.19 

20.32 

1.87+ 

9.0+ 

2.6 

50 

(13) 

6 

*  I  32.48 

30.48 

2.00+ 

7.0+ 

2.0 

60 

dilutions  the  action  is  retarded,  so  that  at  250  cm3  very  nearly 
the  theoretical  amount  of  acid  is  required  to  produce  the  first 
blue  color,  and  at  300  cm3  an  excess  of  21  per  cent  over  the 
theoretical  amount  must  be  added.  If  the  "  after  separa- 
tion "  of  iodine  is  considered  to  be  a  measure  of  the  excess  of 
iodic  acid,  and  if  its  amount  is  accordingly  applied  as  a 
correction,  it  appears  that  for  all  volumes  below  300  cm3  the 
original  thiosulphate  is  completely  destroyed  when  about  90 
per  cent  of  the  theoretical  amount  of  iodic  acid  has  been  added. 
The  presence  of  potassium  iodide  in  the  system  retards  the 
action,  so  that  at  small  volumes  an  excess  of  about  8  per  cent 
of  iodic  acid  must  be  added  to  completely  destroy  the  thiosul- 
phate and  commence  the  separation  of  iodine.  It  is  obvious 
from  the  preceding  experiments  that  the  reaction  between 
iodic  acid  and  sodium  thiosulphate  is  so  indefinite  in  its  nature, 
and  so  dependent  for  completeness  on  conditions  of  time, 
dilution,  and  mass,  that  its  direct  application  as  a  means  of 
standardizing  solutions  must  remain  impracticable. 

*  HI08  added  to  first  blue  color. 

t  Calculated  by  subtracting  from  the  amount  of  iodic  acid  originally 
introduced,  the  volume  of  thiosulphate  of  equal  strength  required  to  bleach 
the  solution  after  standing  twenty  hours. 


THE  COMBUSTION   OF   ORGANIC  SUBSTANCES 
IN  THE   WET  WAY. 

BY  I.  K.  PHELPS .« 

IN  a  former  paper  f  I  have  shown  that  carbon  dioxide  may 
be  estimated  iodometrically  with  a  fair  degree  of  accuracy. 
Inasmuch  as  this  method  is  not  dependent  upon  the  rate  of 
flow  or  rapidity  of  generation  of  the  carbon  dioxide,  it  seemed 
possible  that  some  advantage  might  follow  its  application  to 
the  determination  of  organic  carbon,  oxidized  by  liquid 
reagents. 

Method  of  Oxidation  by  Potassium  Permanganate. 

The  first  experimental  test  in  this  direction  was  made  with 
oxalic  acid,  which  was  oxidized  according  to  the  well-known 
reaction  of  potassium  permanganate  in  the  presence  of  sulphuric 
acid.  The  apparatus  used  was  the  same  as  that  previously 
described  in  the  iodometric  process,  referred  to  above.  It 
consisted,  in  the  main,  of  an  evolution  flask,  and  an  absorption 
flask,  properly  connected.  As  an  evolution  flask,  a  wide- 
mouthed  flask  of  about  75  cm3  capacity  was  used.  This  was 
closed  by  a  doubly  perforated  rubber  stopper,  carrying  a 
separating  funnel  for  the  introduction  of  liquid  into  the  flask 
and  a  glass  tube  of  0.7  cm.  internal  diameter,  which  was 
expanded  to  a  small  bulb  just  above  the  stopper,  to  carry  off 
the  gas.  This  exit  tube  was  joined  by  means  of  a  rubber 
connector  to  a  tube  which  passed  through  the  rubber  stopper 
of  the  absorption  flask,  which  was  an  ordinary  round-bottom 
flask  of  250  cm3  capacity.  This  tube  ended  in  a  valve  of  the 

*  From  Am.  Jour.  Sci.,  iv,  372. 

t  Am.  Jour.  Sci.,  vol.  ii,  p.  70.    Volume  I,  p.  369. 


COMBUSTION  IN  THE    WET   WAY.  63 

Kreider  pattern,*  which  was  enclosed  in  a  larger  tube,  reaching 
nearly  to  the  bottom  of  the  absorption  flask.  The  second  hole 
of  the  stopper  of  this  absorption  flask,  was  filled  by  a  glass 
tube  closed  by  a  rubber  connector  and  screw  pinch-cock. 

The  barium  hydroxide  solution  for  use  in  the  determination 
of  the  carbon  dioxide  was  prepared  by  filtering  a  cold  saturated 
solution  of  the  commercial  salt  into  a  large  bottle,  which  was 
connected  with  a  self-feeding  burette.  The  solution  was 
standardized  in  the  manner  described  in  my  former  paper  by 
boiling  with  an  excess  of  decinormal  iodine  solution  in  an 
ether  wash  bottle.  The  short  tube  of  the  glass  ground  stopper 
of  the  bottle  was  sealed  to  a  Will  and  Varrentrapp  absorption 
apparatus,  which  was  charged  during  the  operation  with  a 
solution  of  potassium  iodide  to  prevent  the  loss  of  elementary 
iodine  in  the  boiling ;  the  long  tube  of  the  bottle  was  used  as 
an  inlet  tube  and  was  closed  externally  by  a  rubber  cap  during 
the  boiling.  After  cooling,  the  excess  of  iodine  used  was 
determined  by  titration  with  decinormal  arsenious  acid  solution 
and  the  iodine  lost  calculated  on  barium  hydroxide  molecule 
for  molecule. 

Potassium  permanganate  was  prepared  for  use  by  dissolving 
the  commercial  salt  in  water,  and  boiling  this  solution,  made 
acid  with  sulphuric  acid,  until  free  from  carbon  dioxide. 
Water  was  also  prepared  free  from  carbon  dioxide  by  boiling 
distilled  water  until  one-third  had  been  driven  off  hi  steam 
and  was  kept  until  used  in  full-stoppered  flasks. 

For  the  first  determinations  of  carbon,  crystallized  ammo- 
nium oxalate  was  weighed  out  and  introduced  into  the  boiling 
flask  with  10-15  cm3  of  pure  water  and  the  flasks  connected 
as  described  above  with  an  appropriate  amount  of  barium 
hydroxide  solution  (3-5  cm8  in  excess  of  the  amount  required 
to  precipitate  the  carbon  dioxide  to  be  determined)  in  the 
absorption  flask.  The  whole  system  was  then  evacuated 
with  the  water  pump  to  a  pressure  of  200-225  mm.  and  the 
oxalate  solution  in  the  boiling  flask  warmed.  An  excess  of 
potassium  permanganate  solution  was  then  run  in  through 
*  Am.  Jour.  Sci.,  1,  p.  132.  Volume  I,  p.  307. 


64  COMBUSTION  OF  ORGANIC  SUBSTANCES 

the  funnel  tube  and  the  mixture  warmed  again,  when  the 
oxidation  of  the  oxalate  was  shown  by  the  carbon  dioxide 
evolved.  The  carbon  dioxide  was  completely  set  free  by 
the  introduction  of  10  cm3  of  sulphuric  acid  (1  : 4)  and  was 
driven  completely  to  the  absorption  flask  by  boiling  for  five 
minutes.  During  the  passage  of  the  gas  into  the  absorption 
flask,  it  was  shaken  frequently  and  was  kept  cool  by  standing 
in  a  dish  of  water  and  by  pouring  cold  water  over  it  from 
time  to  tune.  If,  during  the  boiling,  any  fears  are  enter- 
tained as  to  the  strength  of  the  vacuum  in  the  flasks,  they 
may  be  easily  allayed  by  opening  momentarily  the  stop-cock 
of  the  funnel  tube  and  noting  the  direction  of  the  flow  of 
water,  contained  in  the  funnel.  After  the  boiling  was  ended, 
the  atmospheric  pressure  was  restored  by  allowing  air,  purified 
from  carbon  dioxide  by  passage  through  potash  bulbs,  to 
enter  through  the  funnel  tube  of  the  boiling  flask.  Then 
the  flasks  were  disconnected  and  the  stopper  of  the  absorp- 
tion flask  with  its  attachments  was  removed,  the  valve  and 
its  tube  being  carefully  washed  free  from  barium  hydroxide. 
A  second  stopper,  which  was  provided  with  a  separating 
funnel,  and  a  Will  and  Varrentrapp  absorption  apparatus, 
containing  water  to  serve  as  a  trap,  was  inserted  into  the 
mouth  of  the  absorption  flask  and  the  emulsion  brought  to 
the  boiling  point.  Decinormal  iodine  solution  was  then  run 
in  through  the  funnel  tube  in  sufficient  quantity  to  destroy 
the  larger  part  of  the  excess  of  barium  hydroxide  and  the 
emulsion  brought  to  the  boiling  point  again,  after  which  iodine 
was  again  run  hi  but  this  time  to  the  permanent  red  color  of 
the  excess  of  free  iodine.  After  cooling,  this  excess  of  iodine 
was  determined  by  titration  with  decinormal  arsenious  acid 
solution.  Thus,  the  excess  of  barium  hydroxide  taken  being 
determined  by  the  iodine  lost,  the  barium  hydroxide  used, 
now  hi  the  form  of  carbonate,  was  known,  from  which  the 
carbon  dioxide  which  precipitated  this  carbonate,  may  be 
calculated. 

The  following  results  were  obtained  by  this  procedure. 


IN  THE    WET   WAY. 
TABLE  I. 


65 


Exp. 

Ammonium 

oxalate 
taken. 

Ba02H, 

taken. 

BaO2H2 
found. 

found. 

C02 
calculated. 

Error  on  C02. 

grm. 

grm. 

grm. 

grm. 

grm. 

grm. 

(1) 

0.2522 

0.7267 

0.1170 

0.1565 

0.1561 

0.0004+ 

2 

0.2542 

0.7267 

0.1113 

0.1579 

0.1574 

0.0005+ 

3) 

0.5020 

1.4535 

0.2417 

0.3110 

0.3108 

0.0002+ 

(4) 

0.5058 

1.3954 

0.1753 

0.3131 

0.3131 

O.OOOOi 

(5) 

1.0033 

2.6163 

0.1955 

0.6213 

0.6211 

0.0002+ 

(6) 

1.0003 

2.6951 

0.1836 

0.6189 

0.6192 

0.0003- 

(7) 

1.0010 

2.6163 

0.2037 

0.6192 

0.6197 

0.0005- 

In  experiments  (5)  and  (6),  a  few  drops  of  ammonia 
were  added  to  the  oxalate  solution  before  running  in  the 
permanganate ;  in  (3)  and  (7),  the  permanganate  was  treated 
to  alkalinity  with  barium  hydroxide ;  in  the  remaining  ex- 
periments, (1),  (2),  and  (4),  the  permanganate  was  slightly 
acid  with  the  sulphuric  acid  used  in  its  purification  from 
carbon  dioxide,  as  already  described.  The  results  obtained 
are  good  and  it  is  plain  that  the  oxidation  proceeded  regularly, 
whether  the  first  action  of  the  permanganate  was  in  the 
alkaline  or  slightly  acid  solution. 

Jones  *  has  shown  that  formates  may  be  determined  volu- 
metrically  by  titration  with  potassium  permanganate  in 
alkaline  solution.  In  an  attempt  to  determine  formates  by 
the  process  outlined  above,  the  pure  barium  salt  was  used. 
This  was  prepared  by  treating  the  aqueous  solution  of  formic 
acid  with  pure  barium  carbonate  to  neutrality  and  crystalliz- 
ing the  product.  It  was  proved  pure  by  ignition  and  weigh- 
ing in  the  form  of  carbonate. 

In  making  determinations  of  carbon  in  this  formate,  weighed 
portions  were  introduced  into  the  boiling  flask,  together 
with  sodium  hydroxide  solution,  which  was  taken  in  such 
quantity  as  to  more  than  neutralize  the  acid  in  the  potassium 
permanganate.  Naturally,  the  sodium  hydroxide  must  be 
freed  from  carbonate  —  and  this  was  effected  by  treatment 
with  an  excess  of  barium  hydroxide  and  filtering.  An  excess 


VOL.   II.  —  5 


*  Amer.  Chem.  Jour.,  xvii,  539. 


66 


COMBUSTION  OF  ORGANIC  SUBSTANCES 


of  potassium  permanganate  is  then  run  into  the  flask  and 
the  solution  heated  to  boiling.  An  excess  of  dilute  sulphuric 
acid  is  introduced  into  the  mixture  and  the  carbon  dioxide, 
thus  set  free,  completely  driven  over  to  the  absorption  flask 
and  determined  as  before.  Table  II  shows  results  obtained 
by  the  process. 

TABLE  II. 


Exp. 

Barium 
formate 
taken. 

BaO,H2 
taken. 

BaO,H, 

found. 

CO. 
found. 

CO, 

calculated. 

Error  on  COS. 

grm. 

gnu. 

grin. 

grm. 

grm. 

grm. 

ft 

(3) 

0.5001 
0.5033 
1.0002 

0.9302 
0.9012 
1.6861 

0.1745 
0.1402 
0.1793 

0.1939 
0.1953 

0.3867 

0.1935 
0.1947 
0.3870 

0.0004+ 
0.0006+ 
0.0003- 

(4) 

1.0059 

1.6279 

0.1093 

0.3897 

0.3892 

0.0005+ 

(5) 

1.3750 

2.2529 

0.1820 

0.5315 

0.5320 

0.0005- 

(6) 

1.5028 

2.4419 

0.1754 

0.5816 

0.5814 

0.0002+ 

These  results  show  plainly  that  the  carbon  of  formic  acid 
may  be  determined  accurately  by  the  method  outlined. 

It  was  found  incidentally  that  ammonia  cannot  take  the 
place  of  the  sodium  hydroxide  in  this  process,  probably 
because  the  ammonia  volatilizes  to  the  absorption  flask  dur- 
ing the  boiling  and  is  acted  on  by  the  iodine  subsequently 
used  and  is  thus  registered  as  barium  hydroxide. 

It  is  a  well-known  fact  that  tartrates  are  oxidized  by 
permanganates.  I  have  found,  however,  that  when  tartaric 
acid  is  treated  under  the  conditions  of  analysis  outlined 
above  in  acid  solution,  the  oxidation  is  incomplete;  but  that 
oxidation  is  complete  if  the  tartrate  is  heated  in  a  solution 
alkaline  with  sodium  hydroxide  and  then  acidified  with 
sulphuric  acid. 

The  tartrate  used  was  a  recrystallized  tartar  emetic,  dried 
at  100°  C.  The  following  results  were  obtained  with  such 
a  tartrate  by  this  process. 

It  seems  possible  to  draw  the  general  conclusion  from  the 
results  recorded  that  organic  substances  which  are  oxidized 
completely  by  the  permanganate  may  be  determined  by  the 
process  outlined  above.  It  will  also  be  seen  that  the  use  of 


IN  THE    WET   WAY. 
TABLE  IH. 


67 


Exp. 

Tartar 
emetic, 
taken. 

BaOjH, 

taken. 

BaO^ 

found. 

CO, 
found. 

calculated. 

Error  on  C02. 

grlil. 

grm. 

gnu. 

grm. 

grm. 

grm, 

(1) 

0.5051 

1.2450 

0.1709 

0.2756 

0.2751 

0.0005+ 

(2) 

0.5030 

1.2226 

0.1536 

0.2743 

0.2739 

0.0004+ 

(3) 

0.7509 

1.7355 

0.1401 

0.4094 

0.4091 

0.000:3+ 

(4) 

0.7541 

1.7430 

0.1410 

0.4111 

0.4107 

0.0004+ 

(5) 

1.0018 

2.3456 

0.2187 

0.5458 

0.5456 

0.0002+ 

(6) 

1.0005 

2.2435 

0.1196 

0.5451 

0.5450 

0.0001+ 

the  rubber  stopper  in  the  boiling  flask,  with  due  care  to  prevent 
its  contact  with  the  solution,  does  not  introduce  an  appreciable 
error. 

Wanklyn  and  Cooper  *  and  others  have  noted  the  fact  that 
potassium  permanganate,  whether  in  acid  or  alkaline  solution, 
will  not  oxidize  all  organic  substances  (acetates,  for  example), 
even  at  the  boiling  temperature.  It  is  well  known  that  a 
mixture  of  concentrated  sulphuric  and  chromic  acids  has 
a  much  wider  field  of  action  in  oxidizing  organic  compounds 
than  the  permanganate.  With  hopes  of  applying  this  reagent 
more  widely  to  the  determination  of  organic  carbon,  the 
experiments  about  to  be  recorded  were  tried. 

Method  of  Oxidation  by  Chromic  Acid. 

A  concentrated  mixture  of  chromic  and  sulphuric  acids, 
although  a  much  more  powerful  oxidizer  than  potassium 
permanganate  in  aqueous  solutions,  fails  to  oxidize  completely 
many  organic  compounds.  Thus  Cross  and  Higgin  f  have 
shown  that  carbohydrates  are  among  the  number  of  such 
organic  substances ;  and  later  Cross  and  Be  van  find  that  car- 
bohydrates and  many  other  substances  are  oxidized  to  a  mix- 
ture of  carbon  dioxide  and  monoxide.  Messinger  J  has  proved 
that  carbon  may  be  determined  in  organic  compounds  by  pass- 
ing the  mixed  products,  resulting  from  the  oxidation  with 
chromic  and  sulphuric  acids,  through  a  short  combustion  tube, 

*  Phil.  Mag.  (5),  vii,  138.  t  Jour.  Chem.  Soc.,  xli,  113. 

J  Ber.  Dtsch.  chem.  Ges.,  xxiii,  2756. 


68 


COMBUSTION  OF  ORGANIC  SUBSTANCES 


filled  with  granular  copper  oxide  and  heated  in  a  furnace — all 
of  which  facts  have  been  confirmed  in  my  own  experience. 

Ludwig  *  has  observed  that  the  contact  of  carbon  monoxide 
with  a  mixture  of  chromic  and  sulphuric  acids,  especially 
when  hot,  results  hi  the  oxidation  of  that  gas  to  carbon  dioxide. 
This  fact  suggested  the  idea  of  substituting  for  the  apparatus 
described  above  a  new  form,  adapted  to  retain  the  first  products 
of  oxidation  in  prolonged  contact  with  the  oxidizing  mixture. 

This  apparatus,  shown  in  the  ac- 
companying figure,  by  means  of 
which,  as  the  sequel  shows,  it  has 
been  found  possible  to  extend  the 
availability  of  the  oxidizing  mix- 
ture, is  put  together  as  follows: 
A  thick-walled,  round-bottomed 
flask  of  a  liter's  capacity,  serving 
as  an  oxidizing  chamber,  is  closed 
by  a  rubber  stopper  with  two 
perforations,  through  one  of  which 
passes  the  tube  of  a  separating 
funnel  of  about  100  cm3  capacity. 
The  tube  of  this  funnel  reaches 
nearly  to  the  bottom  of  the  flask 
and  is  drawn  out  at  the  lower  end. 

A  disc  of  platinum  foil  is  hung  in  the  neck  of  the  flask,  nearly 
closing  it,  and  held  in  place  by  a  platinum  wire  passing 
through  the  foil  and  tucked  under  the  rubber  stopper  where 
the  funnel  tube  enters.  The  second  hole  of  the  stopper  is 
filled  by  the  exit  tube,  a  glass  tube  of  0.7  cm.  internal 
diameter.  This  tube  is  expanded  just  above  the  stopper  to 
a  small  bulb,  which  serves  to  prevent  mechanical  loss  of 
the  solid  contents  of  the  flask  during  the  boiling,  and  is 
joined  by  means  of  a  rubber  connector  (provided  with  a  screw 
pinch-cock)  to  the  inlet  tube  of  the  absorption  flask,  which 
is  an  ordinary  500  cm3  round-bottomed  flask.  This  flask  is 
also  closed  by  a  rubber  stopper  with  two  perforations,  through 

*  Ann.  Chem.  (Liebig),  clxii,  47. 


FIG.  19. 


IN  THE   WET  WAY.  69 

one  of  which  passes  the  inlet  tube  described  above  and  through 
the  other  the  exit  tube,  which  is  also  enlarged  to  a  small  bulb 
just  above  the  stopper  and  is  closed  by  a  rubber  connector  and 
screw  pinch-cock.  The  glass  ground  stopper  of  the  funnel 
tube  is  carefully  cleaned  and  lubricated  with  a  thick  solution 
of  metaphosphoric  acid. 

Instead  of  getting  the  vacuum  by  the  water  pump,  it  may 
be  got  almost  as  quickly  and  certainly  more  simply  by  boiling 
the  water  in  the  evolution  flask  and  the  barium  hydroxide 
solution  in  the  absorption  flask  at  the  same  tune  —  both  flasks 
being  connected,  ready  for  making  a  determination.  When 
steam  issues  in  good  quantity  from  the  exit  tube,  the  burner 
is  removed  from  under  the  evolution  flask,  the  attached  pinch- 
cock  closed,  the  burner  under  the  absorption  flask  taken  away, 
and  the  screw  pinch-cock  upon  the  exit  tube  quickly  closed. 
The  flasks  are  then  allowed  to  cool. 

In  making  a  determination,  the  organic  substance  is  weighed 
out  in  a  counterbalanced  bulb,  so  thin  that  it  may  be  easily 
broken  later  and  made  with  a  wide  mouth  for  convenience 
in  introducing  the  solid  substance.  After  the  substance  is 
weighed,  the  mouth  of  the  bulb  is  sealed  by  heating  hi  a 
small  blow-pipe  flame  and  the  tube  introduced  into  the  evolu- 
tion flask,  together  with  an  amount  of  pure  potassium  dichro- 
mate,  which  is  known  to  be  hi  excess  of  that  required  to 
oxidize  the  organic  substance.  The  flasks  are  connected,  as 
already  described,  with  an  appropriate  amount  of  barium 
hydroxide  solution  in  the  absorption  flask  and  10  cm3  of  pure 
water  in  the  evolution  flask,  and  the  vacuum  is  obtained  (as 
described  above)  by  boiling  both  flasks,  the  boiling  being 
stopped  when  the  water  in  the  evolution  flask  has  decreased  to 
2  or  3  cm3.  Naturally,  this  boiling  must  be  so  regulated  as 
not  to  allow  loss  of  the  solid  material  in  either  flask.  The 
vacuum  obtained,  the  tube  containing  the  organic  substance  is 
broken  by  shaking  the  flask,  and  20  cm3  of  concentrated 
sulphuric  acid,  previously  purified  from  organic  material  by 
heating  to  the  fuming  point  with  a  few  crystals  of  potassium 
dichromate,  are  run  in  through  the  funnel  tube,  when  reduc- 


70 


COMBUSTION  OF  ORGANIC  SUBSTANCES 


tion  of  the  chromic  acid  soon  becomes  evident.  While  still 
hot,  the  acid  is  shaken  in  the  flask  violently,  the  platinum  foil 
hung  in  the  neck  serving  to  protect  the  rubber  stopper.  The 
flask  is  warmed  to  approximately  105°  C.,  the  highest  temper- 
ature to  which,  as  shown  by  Cross  and  Be  van,*  a  mixture  of 
chromic  and  sulphuric  acids  may  be  safely  heated  without  the 
disengagement  of  oxygen  gas.  Water  is  then  run  in  until 
the  crystals  of  chromic  anhydride  have  disappeared  and  the 
danger  of  the  evolution  of  oxygen  is  past.  The  solution  is 
heated  to  its  boiling  point,  care  being  taken  that  it  shall  not 
get  under  pressure,  which  can  easily  be  observed  by  opening 
momentarily  the  stop-cock  of  the  funnel  tube  and  noting  the 
direction  of  the  flow  of  water,  contained  in  the  funnel.  The 
flask  is  shaken  and  heated  alternately  for  five  minutes  —  a 
period  of  tune  which  appears  to  be  sufficient  to  bring  about 
the  oxidation  of  the  small  amount  of  carbon  monoxide  origi- 
nally produced.  Then  more  water  (60-70  cm3)  is  introduced 
through  the  funnel,  and  the  stop-cock  between  the  boiling  and 
absorption  flasks  is  opened,  when  the  carbon  dioxide  enters  the 

TABLE  IV. 


Exp. 

Substance 
taken. 

Ba02Hj 
taken. 

Ba02H, 

found. 

C02 

found. 

CO, 

calculated. 

Error  on  C02. 

ANALYSIS  OF  AMMONIUM  OXALATB. 

(1) 
(2) 
(3) 
(4) 
(5) 

grm. 
0.5009 
0.5006 
0.5005 
1.0002 
1.0010 

grm. 
1.3534 
1.3400 
1.3400 
2.5460 
2.5192 

grm. 
0.1469 
0.1308 
0.1343 
0.1347 
0.1094 

grin. 

0.3097 
0.3103 
0.3094 
0.6188 
0.6185 

grm. 

0.3101 
0.3099 
0.3098 
0.6192 
0.6197 

£1*111. 

0.0004- 
0.0004+ 
0.0004- 
0.0004- 
0.0012- 

ANALYSIS  OF  CANE  SUGAR. 

Pj 

% 

0.2001 
0.2000 
0.2001 
0.2014 

1.3926 
1.3926 
1.3926 
1.3400 

0.1905 
0.1936 
0.1857 
0.1279 

0.3085 
0.3077 
0.3097 
0.3111 

0.3088 
0.3086 
0.3088 
0.3108 

0.0003- 
0.0009- 
0.0009+ 
0.0003-f- 

absorption  flask,  which  is  kept  cool  and  shaken  as  before. 

The  contents  of  the  evolution  flask  are  then  heated  to  boiling 

*  Jour.  Chem.  Soc.,  liii,  889. 


IN  THE   WET  WAY.  71 

and  a  slow  current  of  air,  freed  from  carbon  dioxide  by  passing 
through  potash  bulbs,  is  allowed  to  enter  through  the  funnel 
tube  to  keep  the  liquid  from  undue  bumping.  The  boiling  is 
continued  for  fifteen  minutes,  after  which  the  excess  of  barium 
hydroxide  is  determined  iodometrically  and  thus  the  carbon 
dioxide  present  estimated  as  before.  Table  IV  shows  results 
obtained  by  the  treatment  of  crystallized  ammonium  oxalate 
and  cane  sugar,  recrystallized  from  dilute  alcoholic  solution, 
in  this  manner. 

The  results  are  evidently  very  satisfactory. 

The   Determination   of  the   Oxygen  required  to    Oxidize   an 
Organic  /Substance. 

Several  different  methods  have  been  proposed  for  estimating 
the  oxygen  present  in  organic  substances,  depending,  in  gen- 
eral, upon  the  determination  of  the  oxygen  which  must  be 
supplied  to  burn  the  substance  to  a  known  amount  of  carbon 
dioxide  and  water — thus  discovering  by  difference  the  oxygen 
originally  contained  in  the  substance.  Lavoisier  is  said  to  have 
measured  directly  the  oxygen  used  in  burning  organic  sub- 
stances ;  Gay-Lussac  and  The'nard  determined  the  oxygen  used 
by  measuring  the  amount  of  potassium  chlorate  reduced  hi 
burning  the  organic  compound;  Baumhauer*  determined  the 
oxygen  used  by  measuring  the  volume  of  oxygen  entering  the 
combustion  furnace  and  subtracting  the  measure  of  the  gas 
coming  from  the  combustion  tube,  which  was  set  up  according 
to  the  well  known  method  for  determining  carbon  and  hydro- 
gen ;  Stromeyerf  determined  the  amount  of  copper  reduced  by 
the  ignition  of  the  substance  in  copper  oxide;  LadenburgJ 
oxidized  the  substance  by  heating  in  a  sealed  tube  with  a 
known  amount  of  iodic  acid,  determining  at  the  end  of  the 
operation  the  amount  of  iodic  acid  left;  Mitscherlich§  has 
estimated  the  oxygen  in  organic  substances  directly  by  decom- 
posing the  substance  by  ignition  hi  a  stream  of  chlorine  gas, 

/ 

*  Ann.  Chem.  (Liebig),  xc,  228.          t  Ann.  Chera.  (Liebig),  cxvii,  247. 
\  Ann.  Chem.  (Liebig),  cxxxv,  1.        §  Ann.  Phys.  ccvi,  536  (1867). 


72  COMBUSTION  OF  ORGANIC  SUBSTANCES 

estimating  the  oxygen  content  by  determining  the  resulting 
carbon  dioxide  and  monoxide. 

As  it  has  been  shown  in  the  work  described  that  carbon  may 
be  determined  in  organic  substances  by  oxidation  with  chromic 
and  sulphuric  acids  without  the  evolution  of  oxygen  gas,  it 
would  seem  that  the  determination  of  the  oxygen  in  the  sub- 
stance might  be  effected  by  determining  the  amount  of  chromic 
acid  used  in  the  operation,  taking  into  consideration  the 
products  of  combustion.  This  can  be  readily  accomplished  by 
taking  a  weighed  amount  of  pure  potassium  dichromate  as  the 
oxidizing  agent  and  determining,  at  the  end  of  the  operation, 
by  treatment  of  the  residue  with  hydrochloric  acid,  absorp- 
tion of  the  chlorine  evolved  in  an  alkaline  arsenite  of  known 
strength,  and  titration  of  the  excess  of  that  substance  with 
decinormal  iodine  solution,  the  amount  of  chromic  acid  left. 

To  test  the  accuracy  of  the  determination  of  chromic  acid 
under  these  conditions  of  analysis,  weighed  portions  of  pure 
fused  potassium  dichromate  were  introduced  into  a  Voit  flask, 
whose  outlet  tube  was  sealed  to  the  inlet  tube  of  a  Drexel 
wash  bottle,  the  outlet  of  which,  in  turn,  was  sealed  to  a 
Will  and  Varrentrapp  absorption  apparatus.  An  amount  of 
hydrochloric  acid,  more  than  enough  to  completely  reduce 
the  chromate  (15-40  cm3  of  the  strongest  acid),  was  added 
with  20  cm3  of  strong  sulphuric  acid  and  the  total  volume 
made  up  to  120-140  cm3  of  liquid.  The  sulphuric  acid  used 
here  was  purified  from  carbonaceous  matter  (as  in  the  carbon 
determination  above)  by  heating  with  a  few  crystals  of  potas- 
sium dichromate,  the  excess  of  which  was  destroyed  by  hold- 
ing the  acid  at  the  fuming  point  for  about  two  hours,  when 
a  portion  diluted  with  water  gave  no  color  with  potassium 
iodide  and  starch  paste.  Pure  arsenious  oxide,  in  amount 
slightly  in  excess  of  that  required  to  take  up  the  oxygen  to 
be  given  up  by  the  chromate,  was  dissolved  by  the  aid  of 
heat  in  a  solution  of  pure  sodium  hydroxide,  taken  in  such 
quantity  as  to  more  than  neutralize  the  arsenious  acid  and 
the  hydrochloric  acid  used  to  reduce  the  chromate,  and  this 
solution  was  introduced  into  the  Drexel  wash  bottle.  The 


IN  THE    WET  WAY. 


73 


flask  was  then  connected  with  the  wash  bottle,  using  a  thick 
solution  of  metaphosphoric  acid  to  lute  the  joint  between  the 
flask  and  its  stopper.  The  absorption  apparatus  was  charged 
with  a  dilute  solution  of  sodium  hydroxide.  Carbon  dioxide 
was  generated  in  a  Kipp  generator  by  the  action  of  hydro- 
chloric acid  on  marble  and  purified  from  reducing  matter  by 
bubbling  through  a  strong  solution  of  iodine  in  potassium 
iodide  and  finally  washed  with  a  solution  of  potassium  iodide 
alone.  A  slow  stream  of  this  purified  carbon  dioxide  was 
allowed  to  enter  the  inlet  tube  of  the  Voit  flask,  the  con- 
tents of  which  were  then  boiled.  When  concentration  to  a 
volume  of  30-40  cm3  was  reached,  the  boiling  was  discon- 
tinued and,  after  cooling  and  disconnecting  the  flask,  the 
contents  of  the  receiver  were  made  acid  with  sulphuric  acid 
and  then  alkaline  with  acid  potassium  carbonate,  and  the 
excess  of  arsenite  was  determined  by  titration  with  deci- 
normal  iodine  solution.  Sometimes  during  the  reduction  of 
the  chromic  acid,  the  red  fumes  of  the  chlorochromic  an- 
hydride Volatilized  to  the  receiver;  but  since  the  chromic 
acid  thus  produced  is  reduced  later  by  the  arsenite,*  this 
transfer  is  of  no  account  in  the  working  of  the  process. 
The  following  results  were  thus  obtained. 

TABLE  V. 


Exp. 

KjCr2O7 
taken. 

As203 
taken. 

found. 

K2Cr207 
found. 

Error  on 
K2Cra07. 

gnu. 

grill. 

grin. 

grm. 

grm. 

(1) 

5.0002 

5.1025 

0.1144 

4.9447 

0.0555- 

(2) 

5.0018 

5.0799 

0.0526 

4.9849 

0.0169- 

(3) 

5.0005 

5.0801 

0.0582 

4.9782 

0.0223- 

U) 

5.0013 

5.0706 

0.0908 

4.9365 

0.0648- 

The  cause  of  the  error  shown  in  these  experiments  was 
traced  finally  to  too  great  concentration  of  the  sulphuric  acid 
in  the  process.  When  the  boiling  begins  the  chromate  is 
reduced  gradually  and  if  the  evaporation  of  the  water  is 
pushed  too  rapidly,  the  sulphuric  acid  may  reach  a  strength 

*  Browning  Am.  Jour.  Sci.,  i,  35.    Volume  I,  p.  344 . 


74 


COMBUSTION  OF  ORGANIC  SUBSTANCES 


at  which  it  begins  to  cause  the  reduction  of  the  chromic  acid 
with  the  evolution  of  oxygen  instead  of  chlorine. 

The  obvious  remedy  is  to  conduct  the  boiling  operation 
more  slowly.  It  was  found  that,  if  from  5-6  hours'  time  was 
taken  for  the  proper  concentration  of  the  contents  of  the 
Voit  flask,  the  presence  of  the  sulphuric  acid  worked  no 
harm,  as  will  be  seen  from  the  following  results.  Experi- 
ments (1)  and  (5)  were  made  with  5  cm3  of  sulphuric  acid 
present  and  the  others  with  20  cm3,  as  used  before. 

TABLE  VI. 


Bxp. 

K2Cr20T 
taken. 

As20s 
taken. 

As203 
found. 

K*Cr£7 

found. 

Error  on 
K2Cr20,- 

grm. 

grm. 

grin. 

grm. 

grm. 

(1) 

1.0004 

1.0500 

0.0398 

1.0014 

0.0010+ 

(2) 

1.0007 

1.0531 

0.0437 

1.0006 

0.0001- 

(3) 

2.0013 

2.0501 

0.0299 

2.0026 

0.0013+ 

4 

2.0037 

2.0727 

0.0502 

2.0049 

0.0012+ 

(6) 

5.0020 

5.1002 

0.0495 

5.0068 

0.0048+ 

(6) 

5.0037 

6.1018 

0.0513 

5.0066 

0.0029+ 

In  applying  this  method  to  the  determination  of  oxygen 
used  in  the  oxidation  of  an  organic  substance,  the  carbon 
determination  was  made  as  already  described,  the  amount  of 
water  used  being  such  as  to  leave  60-80  cm3  of  liquid  in  the 
boiling  flask  after  the  carbon  dioxide  had  been  driven  to  the 
absorption  flask  by  boiling.  This  liquid  was  then  washed 
into  the  Voit  flask  and  the  chromic  acid  remaining  determined 
by  a  second  distillation  (this  time  with  hydrochloric  acid) 
in  the  manner  described  above.  In  each  of  the  experiments 
recorded  below,  20  cm8  of  purified  sulphuric  acid  were  used 
in  the  carbon  determination  and  35  cm3  of  hydrochloric  acid 
(sp.  gr.  1.2)  in  the  chromic  acid  determination.  The  ammo- 
nium oxalate  used  was  the  pure  crystallized  salt ;  the  phthalic 
acid  was  recrystallized  from  its  water  solution  and  dried  for 
a  short  time  over  sulphuric  acid ;  the  cane  sugar  was  selected 
crystals  of  rock  candy,  recrystallized  from  dilute  alcoholic 
solution  and  dried  for  a  long  time  over  sulphuric  acid;  the 


IN  THE   WET  WAY. 


75 


paper  was  ashless  filter  paper,  dried  to  a  constant  weight  over 
sulphuric  acid;  the  tartar  emetic  was  recrystallized  from 
water  solution  and  air  dried;  the  barium  formate  was  pre- 
pared by  treating  formic  acid  with  an  excess  of  pure  bar- 
ium carbonate,  filtering  hot  and  allowing  the  product  to 
crystallize. 

TABLE   VII. 


Exp. 

Sub- 

staiice 
taken. 

found. 

Error 
on 
CO, 

K2Cr207 
taken. 

A-,0, 

taken. 

found. 

Oxygen 

used. 

Oxygen 
required 
by 
theory. 

Error 
on 
Oxygen. 

ANALYSIS  OF  AMMONIUM  OXALATE. 

(1) 
(2) 

grin. 

1.0122 
1.0019 

grin. 

0.8265 
0.6212 

grm. 
0.0001- 
0.0010+ 

grm. 
2.0009 
2.0002 

grin. 

1.3002 
1.3517 

grm. 
0.0000 
0.0440 

grm. 

0.1160 
0.1147 

grm. 
0.1139 
0.1128 

grm. 
0.0021+ 
0.0019+ 

ANALYSIS  OF  PHTHALIC  ACID. 

(1) 
(2) 

0.1002 
0.1093 

0.2138 
0.2324 

0.0014+ 
0.0007+ 

2.0012 
2.0000 

1.2004 
1.1031 

0.0814 
0.0634 

0.1456 
0.1582 

0.1448 
0.1580 

0.0008+ 
0.0002+ 

ANALYSIS  OF  CANE  SUGAR. 

(1) 
(2) 

0.2025 
0.4012 

0.3117 
0.6166 

0.0008- 
0.0024- 

3.0000 
5.0000 

1.7002 
2.3022 

0.0796 
0.0366 

0.2275 
0.4495 

0.2273 

0.4502 

0.0002+ 
0.0007- 

ANALYSIS  OF  PAPER. 

(1) 

(2) 

0.3034 
0.4523 

0.4932 
0.7334 

0.0010- 
0.0033- 

3.5015 
5.0035 

1.4017 
1.8000 

0.0879 
0.0710 

0.3589 
0.5368 

0.3598 
0.5358 

0.0009- 
0.0010+ 

ANALYSIS  OF  TARTAR  EMETIC. 

(1) 
(2) 

0.5057 
1.0099 

0.2671 
0.5321 

0.0009- 
0.0030- 

2.5018 
3.5003 

1.7000 
1.7520 

0.0766 
0.0198 

0.1459 
0.2911 

0.1462 
0.2919 

0.0003- 
0.0008- 

ANALYSIS  OF  BARIUM  FORMATE. 

(1) 
(2) 

1.0079 
1.5014 

0.3906 
0.5814 

0.0006+ 
0.0005+ 

3.0026 
3.0010 

2.2002 
1.8080 

0.0496 
0.0890 

0.1423 
0.2118 

0.1422 
0.2118 

0.0001+ 
0.0000 

From  these  results,  it  will  be  seen  that  the  process  works 
with  accuracy  upon  a  great  variety  of  organic  substances.  It 
was  found  impossible,  however,  to  determine  the  elements 
in  bodies  which  are  at  the  same  time  volatile  and  hard  to 
oxidize ;  for  instance,  ether  oxidizes  easily  to  acetic  acid 
but  difficultly  beyond  that  stage;  although  the  liquid  acid 


76  COMBUSTION  OF  ORGANIC  SUBSTANCES. 

is  oxidized  vigorously  by  chromic  and  sulphuric  acids,  the 
gaseous  acid  is  hardly  attacked  at  the  temperature  used; 
naphthaline  was  also  found  to  be  volatilized,  and  hence  not 
attacked,  to  such  an  extent  as  to  render  its  determination 
by  this  process  valueless. 


XI 


THE   ESTIMATION   OF   MANGANESE  AS   THE 
SULPHATE   AND   AS   THE   OXIDES. 

BY  F.  A.  GOOCH  AND  MARTHA  AUSTIN.* 

THE  estimation  of  manganese  by  the  conversion  of  salts  of 
that  element  with  volatile  acids  to  the  form  of  the  anhydrous 
sulphate  by  the  action  of  an  excess  of  sulphuric  acid,  evapora- 
tion, and  gentle  heating  was  formerly  a  recognized  procedure. 
On  the  authority  of  Rose,f  however,  this  method  was  set  aside 
on  account  of  the  supposed  difficulty  of  removing  the  excess 
of  acid  without  disturbing  the  composition  of  the  normal  salt. 
Thus,  Oesten,  working  under  Rose's  direction,  obtained,  upon 
submitting  the  crystalline  hydrous  sulphate  MnSO4  .  5H2O  to 
heat,  results  which  may  be  summarized  and  compared  with  the 
results  obtained  by  Rose's  sulphide  method  (the  ignition  of 
the  residue  with  sulphur  in  hydrogen)  as  follows : 


MnS04  .  5H20 
taken. 

MnS04 
found. 

Theory. 

Error. 

MnS 
found. 

Theory. 

Error. 

grm. 

grm. 

gnu. 

grm. 

grm. 

grin. 

grm. 

1.659 

1.043J 

1.037 

0.006+ 

0.597 

0.595 

0.002+ 

1.023§ 

0.014- 

1.481 

0.934} 

0.926 

0.008+ 

0.905§ 

0.021- 

0.725H 

0.201- 

1.430 

0.880§ 

0.893 

0.013- 

0.509 

0.512 

0.003- 

The  residues  remaining  after  the  gentle  ignition  of  the  sulphate 
weighed  apparently  several  milligrams  more  than  should  have 
been  the  case  if  the  salt  had  been  reduced  to  the  normal 
anhydrous  sulphate.  At  higher  temperatures  the  sulphate 


*  From  Am.  Jour.  Sci.,  v,  209.  t  Ann.  Phys.,  clxxxvi,  125  (1860). 

t  Ignited  gently.     §  Ignited  at  low  red  heat.     ||  Ignited  at  strong  red  heat. 


78 


ESTIMATION  OF  MANGANESE  AS   THE 


turned  brown  and  lost  altogether  too  much  weight.  A 
comparison  of  the  errors  of  the  process  in  which  the  ignition 
was  at  low  temperature  with  those  of  the  sulphide  process 
would  seem  to  justify  Rose's  rejection  of  the  former  method 
for  the  latter.  Upon  recalculating  these  results,  however, 
using  atomic  weights  in  use  at  present  —  viz. :  Mn  =  55,  S  = 
32.06,  O  =  16  —  it  becomes  plain  that  the  errors  of  the  two 
processes,  as  shown  in  Oesten's  work,  are  not  very  different 
numerically,  though  with  opposite  signs. 


MnS04  .  5H20 
taken. 

MnS04 
found. 

Theory. 

Error. 

MnS 
found. 

Theory. 

Error. 

grin. 

grm. 

grm. 

grm. 

grm. 

grill. 

grm. 

1.659 

1.043 

1.039 

0.004+ 

0.597 

0.599 

0.002- 

1.481 

0.934 

0.928 

0.006+ 

1.430 

0.509 

0.516 

0.007- 

The  most  uncertain  element  in  these  experiments  is  the 
difficulty,  well-recognized  at  present,  of  getting  the  hydrous 
manganous  sulphate,  upon  which  the  experiments  were  made, 
in  a  perfectly  definite  condition  of  hydration. 

Volhard*  subsequently  studied  the  sulphate  process,  and 
showed  that  manganous  sulphate  may  be  dehydrated,  separated 
from  an  excess  of  sulphuric  acid,  and  brought  into  definite 
condition  for  weighing  as  the  anhydrous  salt  by  careful  and 
protracted  heating  over  a  special  device  of  his  own  —  a  ring 
burner  enclosed  in  a  sheet-iron  casing.  Thus,  on  evaporating 
and  dehydrating  a  solution  of  pure  neutral  manganous  sul- 
phate, Volhard  obtained  the  results  recorded  in  the  following 
statement : 

Residue  of  MnS04  left  by  evaporation  and  dehydration  .  0.1635 
"  after  treatment  with  3  drops  of  H2S04  and  heating 

3  hours 0.1635 

«  after  heating  2  hours 0.1638 

"  after  treatment  with  4  drops  of  H2S04  and  heating 

2£  hours 0.1635 

"  after  heating  3  hours 0.1635 

*  Ann.  Chem.  (Liebig),  cxcviii,  328. 


SULPHATE  AND  AS   THE  OXIDES.  79 

Similar  results  were  obtained  on  evaporating  with  sulphuric 
acid  and  igniting  in  like  manner  an  aqueous  solution  of 
manganous  chloride.  Volhard's  recommendation  of  the  method 
has  not  secured  for  it  the  acceptance  which  its  simplicity  and 
exactness  would  seem  to  demand  —  possibly  because  the 
periods  of  ignition  appear  to  be  considerable  and  the  manner 
of  heating  special. 

In  our  own  experiments  with  the  sulphate  process  we  have 
found  that  special  apparatus  is  unnecessary,  that  the  time  of 
treatment  may  be  short,  and  that  the  process  is  in  every  respect 
simple  as  well  as  very  exact.  We  took  for  a  starting  point 
manganous  chloride  prepared  in  the  manner  to  be  detailed. 
An  aqueous  solution  of  the  so-called  pure  manganous  chloride 
of  commerce  was  boiled  with  pure  manganous  carbonate 
(to  throw  out  aluminum,  iron,  and  chromium),  filtered  and 
precipitated  with  ammonium  sulphide.  The  precipitate  thus 
obtained  was  dissolved  in  a  very  slight  excess  of  hydrochloric 
acid  (to  leave  behind  possible  traces  of  nickel,  cobalt,  and 
copper),  the  solution  was  boiled  to  expel  hydrogen  sulphide 
and  precipitated  with  sodium  carbonate.  The  manganous 
carbonate  thus  thrown  down  was  boiled  repeatedly  with 
successive  portions  of  water,  and  washed  until  the  washings 
were  free  from  chloride.  The  greater  part  of  this  purified 
carbonate  was  dissolved  in  the  least  possible  amount  of  pure 
hydrochloric  acid,  the  reserved  portion  of  the  carbonate  was 
added,  the  mixture  was  boiled,  and  the  solution  of  the  purified 
and  neutral  manganous  chloride  was  filtered  from  the  excess 
of  undissolved  carbonate.  Definite  portions  of  this  solution 
were  precipitated  with  silver  nitrate,  and  from  the  weight  of 
the  silver  chloride  thus  obtained  the  amount  of  manganous 
chloride  present  was  calculated.  Portions  of  the  solution  thus 
standardized  were  drawn,  for  our  experiments,  from  a  burette 
into  a  weighed  platinum  crucible,  sulphuric  acid  was  added 
in  amount  more  than  equivalent  to  the  manganese,  the  solution 
was  evaporated  on  the  water-bath  until  the  water  was  removed, 
and  then,  supported  by  means  of  a  porcelain  ring,  or  triangle, 
within  a  larger  porcelain  crucible  used  as  a  radiator  so  that 


80 


ESTIMATION  OF  MANGANESE  AS   THE 


the  bottom  and  walls  of  the  one  were  distant  from  the  bottom 
and  walls  of  the  other  by  an  interval  of  about  1  cm.,  the 
crucible  was  heated  more  strongly.  The  outer  porcelain 
crucible  may  be  heated  over  a  good  Bunsen  flame  to  a  red 
heat  without  risk  of  overheating  the  manganese  sulphate 
within  the  inner  crucible,  and  the  ignition  may  proceed  as 
rapidly  as  is  consistent  with  the  avoidance  of  mechanical  loss 
by  spattering.  The  results  obtained  by  treatment  of  equal 
portions  (50  cm3)  of  the  same  solution  are  given,  together 
with  the  results  of  standardizing  the  solution  by  precipitation 
with  silver  nitrate,  hi  columns  A  of  the  following  table.  In 
the  other  columns  are  given  comparative  results  got  in  the 
treatment  of  equal  portions  of  several  other  solutions  employed 
subsequently  in  other  work. 


MnSO4 

MnSO4 

calculated 

found  by 

Exp. 

from 
AgCl  found 
in  50  cm3  of 

Exp. 

treatment 
of50cm3of 
solution  A 

Exp. 

MnS04  found  by  treatment  of  50  cm8  of 
various  solutions  with  H2SO4. 

solution  A. 

withH2SO4. 

A. 

A. 

B. 

C. 

D. 

E. 

F. 

G. 

grm. 

grm. 

grm. 

grm. 

grm. 

grm. 

grm. 

grm. 

(1) 

0.3518 

(1) 

0.3513 

(1) 

0.3100 

0.3256 

0.3534 

0.3524 

0.3355 

0.5475 

(2) 

0.3512 

(2) 

0.3514 

(2) 

0.3104 

0.3254 

0.3543 

0.3520 

0.3357 

0.5476 

(3) 

0.3518 

(3) 

0.3096 

These  results  show  plainly  that  the  process  of  estimating 
manganese  in  the  form  of  the  anhydrous  sulphate  is  both 
simple  and  accurate. 

The  estimation  of  manganese  as  the  manganese-manganic 
oxide  Mn3O4,  has  been  so  frequently  criticised  unfavorably 
that  the  method  may  be  said  to  have  passed  from  very  general 
use  excepting  in  certain  cases  in  which  the  directness  of  the 
process  is  a  temptation  to  incur  the  risk  of  some  uncertainty. 
The  production  of  the  other  oxides  of  manganese  in  definite 
condition  is  thought  to  be  even  more  uncertain.  Manganese 
dioxide,  MnO2,  begins,  as  Wright  and  Menke  have  shown  *  to 

*  Jour.  Chem.  Soc.,  xxx,  775. 


SULPHATE  AND  AS   THE  OXIDES.  81 

lose  oxygen  at  a  temperature  (about  210°  C.)  to  which  the 
hydrated  oxide  must  be  heated  to  free  it  from  water,  or  very 
nearly  that  at  which  the  nitrate  is  converted  into  the  dioxide  ; 
so  that  the  chance  of  producing  an  undecomposed  dioxide  by 
the  ignition  of  the  hydrated  dioxide  (the  form  in  which  the 
dioxide  generally  appears  in  analytical  processes),  or  of  the 
nitrate,  is  small.  Manganic  oxide,  Mn2O8,  is  produced,  it  is 
said,  from  the  other  oxides  by  ignition  at  a  low  red  heat  under 
the  ordinary  conditions.  The  manganoso-manganic  oxide, 
Mn3O4,  forms,  presumably,  when  an  oxide  of  manganese  is 
submitted,  under  ordinary  atmospheric  conditions,  to  the  high 
heat  of  the  blast>lamp.  If  the  proportion  of  oxygen  in  the 
surrounding  atmosphere  is  reduced  below  the  normal,  the 
conversion  of  Mn2O3  to  Mn3O4  goes  on  very  easily,  as  Dittmar 
has  shown,*  at  a  temperature  between  the  melting  points  of 
silver  and  aluminum,  while  if  the  proportion  of  oxygen  in  the 
surrounding  atmosphere  rises  much  above  the  normal,  the 
reverse  change,  from  Mn3O4  to  Mn2O3  tends  to  take  place  at 
the  same  temperature.  It  is  not  surprising,  in  view  of  these 
phenomena,  that  the  estimation  of  manganese  as  the  oxide 
Mn3O4  should  have  fallen  into  disrepute;  and  yet,  if  the 
condition  most  favorable  to  the  production  of  that  oxide  —  a 
low  proportion  of  oxygen  in  the  surrounding  air  —  can  be 
maintained  during  the  ignition,  it  is  not  impossible  that  the 
indications  of  the  process  might  prove  to  be,  under  the  con- 
ditions, reasonably  accurate.  Now,  this  may  be  exactly  the 
condition  of  affairs  when  the  ignition  takes  place  ordinarily ; 
for,  if  the  products  of  combustion  displace  the  ordinary  air 
about  the  crucible,  the  proportion  of  oxygen  about  the  oxide 
falls  to  a  low  limit.  We  have  made  the  experiment  of 
enclosing  the  ignited  crucible  within  an  inverted  crucible,  so 
that  the  products  of  combustion  should  be  held  immediately 
about  and  above  the  ignited  oxide,  but  our  experience  has 
shown  that  the  object  in  view  is  attained,  apparently,  quite  as 
well  when  the  ignition  is  so  arranged  that  the  crucible  simply 
rests  well  within  the  upper  part  of  the  flame  of  a  strong  Bunsen 

*  Jour.  Chem.  Soc.,  xvii,  294. 
VOL.  ii.  —  6 


82  ESTIMATION  OF  MANGANESE  AS  THE 

burner,  or  blast-lamp,  in  such  manner  that  an  oxidizing  flame 
covers  nearly  the  entire  wall  of  the  crucible. 

In  the  following  experiments  we  have  put  to  the  test  this 
matter  of  getting  definitely  the  different  oxides  of  manganese. 
We  started  with  a  known  amount  of  pure  anhydrous  sulphate, 
prepared  from  the  pure  chloride  in  the  manner  previously 
described.  This  sulphate  was  converted  by  ignition  into  the 
oxide  — presumably  the  oxide  Mn3O4  —  the  containing  crucible 
being  well  within  the  upper  flame  of  a  powerful  burner. 

In  the  next  step,  this  oxide  was  further  oxidized  by 
moistening  it  with  nitric  acid  and  heating  the  residue  gently 
until  the  evolution  of  fumes  ceased,  the  containing  crucible 
being  placed  well  above  a  porcelain  crucible  used  as  a  radiator 
and  heated  so  that  only  the  bottom  showed  a  faint  red  heat. 
In  this  process  the  attempt  was  made  to  arrest  the  ignition 
at  the  point  where  the  anhydrous  dioxide  was  produced.  As 
the  table  shows,  and  as  would  be  expected,  this  attempt  was 
only  occasionally  and  partly  successful. 

The  residue  of  the  last  process  was  then  submitted  to  a 
higher  heat.  The  platinum  crucible  containing  the  oxide  was 
placed  within  and  touching  the  bottom  of  a  larger  porcelain 
crucible  which  was  heated  to  redness.  Under  these  conditions 
the  temperature  should  not  be  too  hot,  and  the  products  of 
combustion  should  naturally  be  thrown  so  far  away  from  the 
oxide  undergoing  ignition  that  circumstances  should  be 
favorable  for  the  formation  of  the  oxide  Mn2O3.  The  event 
proved  that  the  attainment  of  the  exact  condition  corresponding 
to  the  symbol  Mn2O3  is  a  matter  of  some  uncertainty. 

Next,  the  oxide  was  subjected  to  the  highest  heat  of  a  strong 
Bunsen  burner  (or  in  some  cases,  the  broad  flame  of  a  blast 
lamp),  the  crucible  being  well  surrounded  by  the  products  of 
combustion.  The  results  of  this  treatment,  it  will  be  seen, 
agree,  with  a  single  exception  out  of  ten  experiments,  reason- 
ably well  with  the  theory  for  Mn8O4. 

By  treating  the  final  oxide  with  nitric  acid  and  repeating 
the  cycle  of  operations  described,  the  observations  of  the 
phenomena  were  multiplied,  until,  finally,  the  oxide  formed 


SULPHATE  AND  AS   THE   OXIDES. 


83 


I  4 


o'oo 


CO  CO  CO 

ooo 


4  I 


OOO 


II      II 


oo     oo 


li 


rH  t~  <M  OS 
OO  00  GO  GO 
O  OOO 


t—     O  T*     COT* 

O   O  O*   O"  O 


1-1   O  OO   <N  U3 


,4444-     J^^ 

'  C^  O  CO  O        O  T— (  i— I 

'oooo     ooo 


oooo  ooo 


88  888 


oo  ooo 


(M  t- 

CO  CO  CO  CO  CO 

CO  CO  CO  CO  CO 

o'  o  ooo 


CCOOSO        COCDOS 


4  I  4- 


oooooooo 


^H  CO  <M  O  (N  7^  t— 

o  o  o'  o'  o  o  o 


I 


•M  CO  CD  Oi  O  CO  CD 


84  ESTIMATION  OF  MANGANESE,  ETC. 

last  was  treated  with  sulphuric  acid,  ignited  in  the  manner 
previously  detailed,  and  weighed  as  the  anhydrous  sulphate, 
thus  showing  that  no  significant  loss  of  material  had  taken 
place  in  the  series  of  manipulations.  The  table  comprises  the 
results  of  these  experiments.  The  numbers  in  parentheses 
indicate  the  order  of  treatment. 

The  inference  is  plain  that  the  estimation  of  the  manganese 
in  the  form  of  the  manganoso-manganic  oxide,  Mn8O4,  is  by  no 
means  to  be  considered  utterly  untrustworthy  when  the  process 
is  conducted  in  the  manner  described,  though  it  must  be 
recognized  that  an  irregular  result  may  occur  occasionally. 
The  danger  of  accepting  such  an  irregularity  as  a  correct 
indication  may  be  eliminated  to  a  very  considerable  extent  if 
the  precaution  is  taken  invariably  to  moisten  the  ignited  oxide 
with  nitric  acid,  and  ignite  again.  The  indications  of 
harmonious  results  thus  got  may  be  taken  with  a  fair  degree 
of  confidence.  However,  it  is,  in  our  judgment,  by  far  the 
wiser  and  simpler  plan  to  convert  an  oxide  of  manganese 
obtained  in  course  of  analysis  into  the  sulphate  and  to  weigh 
the  manganese  in  that  form. 


XII 

ON  THE  CONDITION   OF  OXIDATION  OF 

MANGANESE    PRECIPITATED  BY  THE 

CHLORATE  PROCESS. 

BY  F.  A.  GOOCH  AND  MARTHA  AUSTIN.* 


HANNAY,|  wno  was  the  fr"8*  ^°  propose  the  precipitation  of 
manganese  from  its  solution  in  nitric  acid  by  the  use  of 
potassium  chlorate,  states  that  precipitation  is  complete,  but 
that  the  oxide  produced  is  not  of  constant  composition. 
While,  therefore,  precipitation  by  this  method  serves  an 
excellent  purpose  in  separating  manganese  from  other  sub- 
stances, it  was  Hannay's  opinion  that  reliance  cannot  be  placed 
upon  the  determination  of  the  oxygen  value  of  the  oxide  to 
estimate  the  manganese.  Beilstein  and  Jawein,f  who  proposed 
the  same  method,  subsequently,  regarded  the  precipitate  as 
the  oxide  MnO2.  Hannay's  reaction  was  developed,  inde- 
pendently, by  Hampe  §  and  Ford||  into  the  method  which  is 
now  known  as  the  "  chlorate  process"  for  the  estimation  of 
manganese.  The  discussion  of  the  exact  condition  of  the 
precipitated  oxide  was  very  active  ten  years  ago,  and  occasional 
echoes  of  it  are  heard  at  the  present  day  ;  and  yet,  in  all  this 
discussion  we  find  no  account  of  an  adequate  test  of  the 
process  upon  an  exact  amount  of  a  salt  of  manganese  known 
to  be  pure.  The  discussion  for  the  most  part  has  centered 
about  the  degree  of  oxidation  of  the  precipitate,  but  there  is 
obviously  another  condition  to  be  taken  into  account,  viz.  : 
the  possibility  of  the  mechanical  inclusion  of  the  comparatively 
insoluble  chlorate  in  the  precipitated  oxide.  As  to  the 
existence  of  the  latter  source  of  error  we  have  had  in  the 

*  From  Am.  Jour.  Sci.,  v,  260.  f  Jour.  Chem.  Soc.,  vol.  xxxiii,  269. 

t  Ber.  Dtsch.  chem.  Ges.,  xii,  1528.        §  Chem.  Central-Blatt,  1885,  714. 
||  Trans,  last.  Am.  Min.  Eng.,  ix,  347. 


86 


CONDITION  OF  OXIDATION  OF  MANGANESE 


course  of  our  work  very  strong  affirmative  evidence,  the 
apparent  condition  of  oxidation  of  the  precipitate  being 
sometimes  so  high  as  to  be  otherwise  inexplicable.  This 
difficulty  does  not  occur,  however,  when  a  more  soluble 
chlorate  is  chosen  to  do  the  work  of  oxidation,  and  we  have 
found  quite  as  convenient  and  much  safer  the  substitution  of 
sodium  chlorate  for  the  comparatively  insoluble  potassium 
chlorate.  Besides,  the  rapidity  with  which  the  sodium 
chlorate  is  decomposed  makes  its  use  an  advantage. 

With  regard  to  the  completeness  of  the  precipitation  our 
experience  has  shown  that  with  due  precaution  the  method  is 
practically  perfect.  Thus,  after  boiling  manganous  nitrate 
(free  from  chlorides  and  sulphates)  with  strong  nitric  acid 
(85  cm3)  and  sodium  chlorate  (5  grm.)  for  five  minutes, 
adding,  subsequently,  15  cm3  of  the  nitric  acid  and  a  few 
crystals  more  of  the  sodium  chlorate,  and  discontinuing  the 
heating  as  soon  as  the  liquid  again  boils,  the  insolubility  of  the 
manganese  is  so  great  that  no  more  than  insignificant  traces 
may  be  recovered  from  the  filtrate  after  cooling,  filtering  upon 
asbestos  and  washing  with  water.  The  test  for  manganese  in 
the  filtrate  and  washings  was  made  after  evaporation  and 
solution  of  the  residue  in  distilled  water  by  treating  the  hot 
solution  with  bromine  and  ammonia.  In  the  first  division  of 
the  table  below  are  results  obtained  by  treating  the  manganese 

TABLE  I. 


MnSO« 

taken. 

Mn  found  by  KI 
treatment  in 
filtrate. 

Mn  found  by 
As2O3  treatment 
in  filtrate. 

grm. 
0.3361 
0.3361 
0.3361 
0.3361 
0.3361 
0.3361 
0.3361 

grm. 

None. 
None. 
0.00006 
0.00005 
0.00002 
0.00008 
None. 

grm. 

0.4128 
0.4128 
0.4128 
0.4128 

'..'.'. 

000003 
000003 
Trace. 
Trace. 

PRECIPITATED  BY  THE   CHLORATE  PROCESS.       87 

precipitated  from  the  filtrate  with  potassium  iodide  and 
sulphuric  acid,  the  iodine  set  free  being  determined  by  sodium 
thiosulphate.  In  the  second  series  the  precipitated  manganese 
dioxide  was  reduced  by  a  known  amount  of  decinormal 
arsenious  acid  and  the  amount  remaining  unoxidized  was 
estimated  by  titration  with  iodine  in  the  presence  of  acid 
potassium  carbonate. 

It  will  be  seen  that  in  no  case  did  the  manganese  which 
escaped  precipitation  —  that  which  corresponded  to  the  iodine 
freed  or  the  arsenious  acid  oxidized  —  exceed  0.0001  grm. 
Plainly  this  modified  method  of  handling  the  chlorate  process 
may  be  trusted  to  precipitate  the  manganese  with  gratifying 
rapidity  and  approximation  to  completeness.  Our  experience 
has  shown  plainly  that  prolonged  boiling  results  in  a  considera- 
ble loss  of  manganese  (from  0.0010  —  0.0030  grm.).  This  we 
think  is  due  to  the  solvent  effect  of  the  lower  oxides  of  nitro- 
gen naturally  produced  (as  is  always  the  case  in  boiling  nitric 
acid)  after  the  chlorine  dioxide  has  been  thoroughly  expelled. 
An  excess  of  the  chlorate  at  the  end  of  the  boiling  seems  to  be 
essential  and  a  slight  yellow  color  in  the  solution,  due  to  chlo- 
rine dioxide,  is  a  favorable  indication  rather  than  the  reverse. 
We  find  it  best  to  filter  the  undiluted  nitric  acid,  under  pressure, 
upon  asbestos  on  a  perforated  cone  with  a  filtering  surface  of 
about  40  cm2.  The  dilution  of  the  nitric  acid  before  filtra- 
tion tends  to  produce  some  solubility  of  the  manganese,  and 
the  loss  then  introduced,  though  trifling  if  the  filtration  is 
rapid,  may  be  considerable  if  the  process  of  filtration  is  pro- 
longed, as  is  the  case  in  the  method  approved  by  the  "  Verein 
der  deutschen  Eisenhiitteleute."  * 

Our  experiments  upon  the  chlorate  process  have  been  made 
with  manganous  chloride  prepared  as  detailed  in  a  former 
paper,  viz.:  by  boiling  manganous  chloride  with  manganous 
carbonate,  precipitating  the  filtered  solution  with  ammonium 
sulphide,  dissolving  the  washed  manganous  sulphide  in  dilute 
hydrochloric  acid,  precipitating  the  solution  thus  obtained  with 
sodium  carbonate  (after  boiling  out  hydrogen  sulphide),  dis- 

*  Von  Reis,  Zeitschr.  angew.  Chem.,  1891,  376. 


88         CONDITION  OF  OXIDATION  OF  MANGANESE 

solving  the  greater  part  of  the  manganous  carbonate  (thor- 
oughly washed  by  repeatedly  boiling  it  in  successive  portions 
of  water)  in  the  least  amount  of  hydrochloric  acid,  and  boiling 
the  solution  thus  obtained  with  the  remainder  of  the  pure 
carbonate  and  filtering.  The  standard  of  the  solution  thus 
prepared,  neutral  and  probably  very  pure,  was  fixed  by  evapo- 
rating definite  portions  with  sulphuric  acid  and  weighing 
the  residue  as  the  normal  sulphate  in  accordance  with  the  pro- 
cedure outlined  in  a  former  paper.* 

Any  method,  by  means  of  which  the  oxidizing  power  of  the 
higher  oxygen  compounds  of  manganese  is  discoverable,  may, 
obviously,  be  employed  to  determine  the  condition  of  the 
manganese  precipitated  in  this  chlorate  process.  Convenient 
processes  for  the  determination  of  the  available  oxygen  in  the 
higher  oxides  of  manganese  are  the  iodometric  methods  of 
Bunsen  and  Pickering.  Bunsen's  method  is  applicable  to  any 
of  the  higher  oxygen  compounds  of  manganese  —  though  some- 
what inconvenient  because  it  involves  the  distillation  of  the 
chlorine  liberated  by  the  action  of  strong  hydrochloric  acid 
upon  the  substance  and  its  collection  in  potassium  iodide,  the 
iodine  thus  set  free  being  estimated  by  standard  thiosulphate. 
According  to  Pickering's!  method  the  higher  oxide  is  treated 
immediately  with  potassium  iodide  and  hydrochloric  acid  and 
the  iodine  liberated  is  estimated  by  sodium  thiosulphate. 
Plainly,  the  latter  method  is  limited  to  the  treatment  of  the 
less  refractory  or  more  finely  comminuted  oxides,  and  it  fails 
in  the  presence  of  ferric  salts  and  all  other  substances  capable 
of  liberating  iodine  from  the  acidified  iodide. 

Still  another  general  iodometric  method  for  determining  the 
oxygen  value  of  the  higher  oxides  of  manganese  is  suggested 
by  Deshayes's  titration  of  permanganic  acid  in  nitric  acid  by 
means  of  standard  arsenious  acid.J  Our  experience  in  follow- 
ing out  this  idea  shows  that  the  precipitated  oxides  of  manga- 
nese, as  well  as  the  soluble  permanganate,  may  be  easily 
reduced  with  the  aid  of  gentle  heating  by  arsenious  acid  in 

*  Am.  Jour.  Sci.,  v,  209.    This  volume,  p.  77. 

t  Jour.  Chem.  Soc.,  xxxvii,  128.        \  Bull.  Soc.  Chim.,  xxix,  541. 


PRECIPITATED  BY  THE  CHLORATE  PROCESS.       89 

the  presence  of  sulphuric  acid,  and  that  the  determination  of 
the  excess  of  arsenious  acid  by  titration  with  iodine  after  neu- 
tralization of  the  free  sulphuric  acid  by  an  alkaline  carbonate 
gives  exact  data  for  estimating  the  oxidizing  power  of  the 
manganese  compound.  We  found,  however,  that  if  the  iodine 
is  allowed  to  come  into  contact  with  the  manganous  carbonate 
thrown  down  by  the  alkaline  carbonate,  as  is  inevitable 
for  at  least  short  intervals  during  the  titration  of  the 
arsenious  acid  in  presence  of  the  floating  carbonate,  the 
danger  arises  of  more  or  less  reoxidation  of  the  manganous 
carbonate  by  the  iodine  and  the  consequent  introduction  of 
error.  Fortunately  the  difficulty  may  be  obviated  by  adding 
to  the  solution,  while  still  acid,  enough  tartaric  acid  or  alkaline 
tartrate  to  prevent  the  precipitation  of  the  manganese  in  the 
subsequent  neutralization  by  the  bicarbonate. 

Confining  our  attention  to  the  last  two  simpler  iodometric 
methods  —  the  reduction  of  the  higher  oxide  by  an  acidified 
iodide  on  the  one  hand  and  by  arsenious  acid  on  the  other  — 
we  made,  first,  some  experiments  to  determine  the  accuracy 
with  which  manganese  may  be  thus  estimated.  We  used 
for  the  manganese  compound  of  known  oxidizing  power  a 
solution  of  potassium  permanganate  filtered  carefully  through 
asbestos  and  standardized  against  ammonium  oxalate  which 
had  been  found  to  be  the  exact  equivalent  of  a  specially 
prepared  lead  oxalate.  For  each  experiment  a  definite  por- 
tion of  this  solution  was  drawn  from  a  burette  and  treated 
with  a  solution  of  pure  manganous  sulphate  until  the  color 
of  the  permanganate  had  vanished,  thus  precipitating  a 
hydrous  oxide  approximating  quite  closely  probably  to  the 
condition  of  oxidation  of  the  dioxide,  but  containing  at  all 
events,  whatever  its  actual  composition  might  be,  the  exact 
amount  of  available  oxygen  originally  in  the  permanganate. 
In  the  experiments  of  the  following  table  this  precipitate 
was  treated  with  a  solution  of  potassium  iodide  (6  grm.)  and 
tartaric  acid  (10  grm.),  by  which  the  freshly  prepared  hydrate 
is  dissolved  quite  as  well  as  by  the  iodide  and  hydrochloric 
acid  of  Pickering's  original  method  and  with  less  risk  of 


90 


CONDITION  OF  OXIDATION  OF  MANGANESE 


evolution  of  iodine  outside  the  main  reaction.  From  the 
iodine  found  by  titration  with  this  sulphate  we  have  cal- 
culated the  weight  of  manganese  dioxide  which  would  liberate 
it;  and  a  comparison  of  this  value  with  the  amount  of  the 
dioxide  theoretically  precipitated  by  the  interaction  of  the 
known  permanganate  and  the  sulphate,  upon  the  assumption 
that  two  molecules  of  the  former  throw  down  five  molecules 
of  the  hydrated  dioxide,  should  disclose  the  error  of  the 
analytical  process  when  applied  to  the  estimation  of  man- 
ganese dioxide.  In  all  probability  the  assumption  that  it 
is  the  dioxide  that  is  precipitated,  and  which  afterwards  acts 
upon  the  iodide,  is  not  quite  true  under  the  conditions,  since 
the  precipitation  takes  place  in  presence  of  an  excess  of  a 
manganous  salt ;  but  for  our  purpose  it  does  not  matter,  since 
we  are  in  effect  simply  dealing  with  the  oxidizing  power  of 
a  known  amount  of  permanganate. 

TABLE  II. 


Mn  theoretically 
precipitated  by 
KMnO4  as  MnO2. 

Mn  in  MnO2 
corresponding 
to  iodine  found. 

Error  of  the  analytical 
process  (applied  to 
MuO2)  in  terms  of  Mn. 

grm. 
0.1351 
0.1351 
0.1351 
0.1351 
0.1351 
0.1351 

gnu. 
0.1347 
0.1347 
0.1350 
0.1353 
0.1358 
0.1353 

grin. 

0.0004- 
0.0004- 
0.0001- 
0.0002+ 
0.0007+ 
0.0002+ 

Thus,  it  is  obvious  that  the  mean  error  of  the  results 
is  practically  inconsiderable,  varying  between  extremes  of 
—  0.0004  grm.  and  +  0.0007  on  0.1351  grm.  of  manganese 
dioxide. 

In  the  experiments  of  Table  III  the  precipitated  oxide  was 
treated  with  an  excess  of  a  standard  arsenious  acid  solution 
and  5  cm3  of  sulphuric  acid  of  half  strength,  and  the  whole 
was  heated  until  the  manganese  dioxide  was  dissolved.  To 
this  liquid  was  added  tartaric  acid  (10  grm.)  to  prevent  the 
precipitation  of  the  manganese  and  the  oxidation  by  iodine  in 


PRECIPITATED  BY  THE   CHLORATE  PROCESS.       91 


the  subsequent  titration,  the  acid  was  neutralized  by  acid 
potassium  carbonate,  and  the  arsenic  still  remaining  in  the 
arsenious  condition  titrated  by  standard  iodine. 


TABLE  IIL 


Mn  precipitated  by 
action  of  KMnO4 
on  MnSO4  as  MnO2. 

MninMnO2 
corresponding  to 
As-jOg  oxidized. 

Error  of  the  process 
in  terms  of  Mn. 

grm. 

grm. 

grm. 

0.1392 

0.1396 

0.0004- 

0.1109 

0.1117 

0.0008- 

0.1112 

0.1117 

0.0005- 

0.1109 

0.1117 

0.0008- 

0.1109 

0.1117 

0.0008- 

0.1117 

0.1125 

0.0008- 

It  is  clear  that  either  of  these  methods  of  reduction,  the 
action  of  an  acidified  iodide  or  that  of  arsenious  acid,  is 
capable  of  yielding  fairly  accurate  indications  when  we  have 
to  deal  with  a  pure  salt  of  manganese.  When,  however,  the 
manganese  is  associated  with  a  considerable  amount  of  iron, 
as  is  frequently  the  case,  it  becomes  a  matter  of  necessity  to 
separate  the  manganese  before  attempting  its  estimation. 
For  this  purpose  the  "  chlorate  process  "  is  by  far  the  simplest 
of  those  generally  used,  and  though  it  has  been  the  subject 
of  much  discussion,  it  is  at  present  the  method  of  separation 
most  widely  used  by  practical  chemists,  whether  the  final 
estimation  of  the  manganese  is  made  gravimetrically  as  in 
Ford's  process,  or  volumetrically,  as  in  the  methods  of 
Volhard,  Williams,  or  Pattinson. 

Definite  portions  of  the  solution  of  pure  manganous  chloride 
were  drawn  from  a  burette  into  an  Erlenmeyer  flask  of  300  cm3 
capacity,  evaporated  to  dryness,  precipitated  by  the  "  chlorate 
process  "  with  the  modifications  given  in  detail  above.  The 
oxide,  after  careful  washing,  was  returned  with  the  asbestos 
to  the  flask  and  treated  by  one  or  other  of  the  methods  to 
be  described.  It  was  either  treated  with  potassium  iodide 
(5  grm.)  and  sulphuric  acid  (10  cm3)  of  half  strength,  the 
iodine  set  free  being  estimated  by  thiosulphate ;  or  it  was 


92 


CONDITION  OF  OXIDATION  OF  MANGANESE 


heated  with  an  excess  of  standard  arsenious  acid  and  10  cm3 
of  sulphuric  acid  of  half  strength,  and  after  cooling,  adding 
5  grm.  of  the  Rochelle  salt  and  neutralizing  with  acid 
potassium  carbonate,  the  arsenious  acid  remaining  unoxidized 
was  estimated  with  standard  iodine.  In  Table  IV  are  given 
the  results  obtained  in  this  work. 

TABLE   IV. 


Mn  taken  in 
the  form  of 
manganous 
chloride. 

Mn  found  upon 
the  hypothesis 
that  MnOi  is 
precipitated. 

Error. 

Mn  found  in  the  nitrate 
after  evaporation  and 
treatment  with  bromine 
and  ammonia. 

BY  REDUCTION  WITH  POTASSIUM  IODIDE. 

grin.. 

0.1225 
0.1225 
0.1225 
0.1225 

grm. 
0.1183 
0.1177 
0.1180 
0.1169 

grill. 

0.0042- 
0.0048- 
0.0045- 
0.0056- 

grm. 
0.00006 
Trace. 
0.00008 
Trace. 

BY  REDUCTION  WITH  ARSENIOUS  ACID. 

0.1222 
0.1222 
0.1222 
0.1222 
0.1222 
0.1222 
0.1222 
0.1222 
0.1222 

0.1189 
0.1191 
0.1199 
0.1200 
0.1186 
0.1187 
0.1189 
0.1194 
0.1205 

0.0033- 
0.0031- 
0.0023- 
0.0022- 
0.0036- 
0.0035- 
0.0033- 
0.0028- 
0.0017- 

Not  determined. 
Not  determined. 
Not  determined. 
Not  determined. 
None. 
0.0001 
0.0002 
Trace. 
0.0001 

The  results  show  plainly  that,  while  the  manganese  is  so 
completely  precipitated  in  the  chlorate  process  of  oxidation 
when  properly  conducted  that  only  insignificant  traces  may 
escape,  the  condition  of  oxidation  cannot  be  taken  to  be 
that  of  the  dioxide.  The  average  error  thus  put  upon  the 
determination  of  the  manganese  known  to  be  present  is  more 
than  2  per  cent.  It  follows,  as  a  matter  of  course,  that  the 
indications  of  any  process  which  rests  upon  the  assumption 
that  the  oxygen  value  of  the  manganese  compound  preci- 
pitated in  the  chlorate  process  corresponds  to  that  of  the 
dioxide  must  of  necessity  be  erroneous.  If,  therefore,  the 
chlorate  method  is  to  be  employed  for  the  separation  of 


PRECIPITATED  BY  THE   CHLORATE  PROCESS.       93 

the  manganese,  it  is  obvious  that  precautions  must  be  taken 
to  secure  a  definite  condition  of  oxidation  of  the  manganese 
before  processes  which  depend  upon  the  oxygen  value  of 
the  higher  oxide  may  be  applied  to  the  estimation  of  that 
element.  The  process  which  in  our  hands  seems  to  give 
the  oxide  in  definite  condition  is  based  upon  the  observations 
of  Wright  and  Menke*  that  a  dilute  solution  of  potassium 
permanganate  acting  in  excess,  at  80°  C.,  hi  the  presence  of 
zinc  sulphate,  and  in  thorough  mixture  upon  manganous 
sulphate,  yields  an  oxide  which,  though  combined  with  alkali, 
holds  the  oxygen  exactly  in  the  proportion  corresponding  to 
the  dioxide.  Three-fifths  of  the  manganese  in  such  a  pre- 
cipitate represents  the  amount  of  that  element  originally 
present  in  the  manganous  salt.  In  the  following  table  are 
given  the  results  of  experiments  in  which  manganese  was 
determined  iodometrically  after  the  interpolation  of  the 
permanganate  treatment. 

In  these  experiments  a  solution  of  manganous  chloride  of 
known  strength  was  drawn  from  a  burette,  evaporated  to 
dryness  in  a  small  beaker,  heated  with  nitric  acid  until  there 
was  no  evidence  of  the  presence  of  nitrogen  oxides.  Strong 
nitric  acid  was  poured  in  until  the  volume  was  85  cm3,  sodium 
chlorate  (5  grm.)  was  added  carefully,  the  liquid  was  boiled 
five  minutes,  more  nitric  acid  (15  cm3)  and  a  few  crystals  of 
the  chlorate  were  introduced,  and  the  solution  brought  to 
boiling  temperature  again.  After  cooling,  the  liquid  was 
filtered  on  asbestos  and  washed  with  water,  and  the  oxide 
upon  the  asbestos  and  walls  of  the  beaker  was  dissolved  in 
2  cm3  of  hydrochloric  acid.  After  diluting  a  little  the 
solution  was  evaporated  with  5  cm8  of  strong  sulphuric  acid 
until  no  more  hydrochloric  acid  remained.  The  solution  of 
manganous  sulphate  (not  exceeding  0.5  grm.  of  the  salt), 
very  nearly  neutralized  by  potassium  carbonate,  was  mixed 
with  a  solution  of  zinc  sulphate  (2  grm.)  and  a  freshly  and 
carefully  filtered  dilute  solution  of  potassium  permanganate 
(1.5  grm.  of  the  salt)  ;  the  liquid,  amounting  now  to  about 
*  Jour.  Chem.  Soc.,xxxvii,  36. 


94 


CONDITION  OF  OXIDATION  OF  MANGANESE 


500  cm3,  was  heated  to  80°  C.,  and  acid  potassium  carbonate 
added,  in  quantity  a  little  more  than  enough  to  neutralize  the 
remnant  of  the  acid  present.  The  precipitate  was  collected 
upon  asbestos,  and  after  careful  washing  was  returned  to  the 
flask  in  which  the  precipitation  had  been  made.  The  oxygen 
value  of  the  oxide  was  determined  by  one  or  other  of  the 
methods  described.  In  the  one  case  the  flask  was  fitted  with 
a  paraffined  stopper  having  two  bores,  one  holding  a  Will 
and  Varrentrapp  absorption  apparatus  (in  which  a  solution  of 
potassium  iodide  dissolved  any  escaping  iodine),  the  other  a 
small  separating  funnel.  Sulphuric  acid  and  potassium  iodide 
in  solution  were  run  in  through  the  funnel,  the  iodine  set 
free  was  titrated  with  thiosulphate  —  the  amount  of  man- 
ganese being  reckoned  from  the  iodine  set  free.  The  results 
of  this  work  follow  in  the  first  part  of  Table  V.  In  the 
second  case  the  dioxide  obtained  in  the  manner  described  above 
was  reduced  by  warming  gently  with  a  decinormal  solution 
of  arsenious  acid.  After  cooling,  and  neutralizing  with  acid 

TABLE  V. 


Mn  taken 
in  the 
form  of  chloride. 

Mn  found  upon 
the  hypothesis  that 
MnO2  is  the  oxide 
finally  obtained. 

Error. 

BY  REDUCTION  WITH  POTASSIUM  IODIDE. 

grm. 

grm. 

grm. 

0.0643 

0.0637 

0.0006- 

0.0643 

0.0642 

0.0001- 

0.0643 

0.0642 

0.0001- 

0.0651 

0.0651 

0.0000 

0.1125 

0.1121 

0.0004- 

0.1125 

0.1121 

0.0004- 

0.1125 

0.1120 

0.0005- 

0.1214 

0.1206 

0.0008- 

0.1214 

0.1207 

0.0007— 

0.1214 

0.1223 

0.0009+ 

0.1214 

0.1214 

0.0000 

BY  REDUCTION  WITH  ARSENIOUS  OXIDE. 

0.1213 

0.1212 

0.0001- 

0.1213 

0.1201 

0.0012- 

0.1213 

0.1203    . 

0.0010- 

0.1213 

0.1208 

0.0005- 

PRECIPITATED   BY  THE   CHLORATE  PROCESS.       95 

potassium  carbonate  in  the  presence  of  Rochelle  salt,  the 
excess  of  the  arsenious  acid  was  estimated  with  iodine  in  the 
presence  of  starch.  The  estimation  by  this  method  gave 
the  results  recorded  hi  the  second  part  of  Table  V. 

These  results  show  plainly,  that  if  the  precautions  to  which 
attention  has  been  directed  are  taken,  viz.:  dilution  of  the 
solution  and  heating  to  80°  C.,  presence  of  zinc  sulphate, 
and  (most  essential  of  all)  the  almost  complete  neutralization 
of  free  acid  before  the  addition  of  the  potassium  permanganate, 
the  manganese  dioxide  precipitated  by  the  chlorate  process 
from  pure  manganous  nitrate  may  subsequently,  after  reduc- 
tion, be  brought  by  the  permanganate  treatment  so  nearly  to 
the  full  degree  of  oxidation  represented  by  the  symbol  MnO2, 
that  the  amount  of  manganese  originally  treated  may  be 
calculated  with  a  very  fair  degree  of  accuracy  from  the 
oxygen  value  of  three-fifths  of  the  oxide  found.  We  do  not 
recommend  this  procedure  as  a  rapid  analytical  method ;  our 
purpose  is  accomplished  when  the  fact  is  brought  plainly  to 
view  that  the  oxide  precipitated  by  the  chlorate  process  is 
not  the  dioxide,  but  that  it  may  be  made  such  by  subsequent 
treatment. 


XIII 

ON  THE  ESTIMATION  OF  MANGANESE 
SEPARATED  AS  THE  CARBONATE. 

BY  MAKTHA  AUSTIN  * 

THE  estimation  of  manganese  precipitated  as  the  manganous 
carbonate,  when  that  salt  is  obtained  by  the  action  of  sodium 
or  potassium  carbonate,  has  been  regarded  as  very  undesirable 
for  the  reasons  that,  even  if  the  conditions  of  the  precipitate 
is  such  that  it  does  not  run  through  the  filter,  the  manganous 
carbonate  can  never  be  freed  entirely  from  alkaline  salt,  and 
that  the  conversion  of  the  carbonate  to  the  manganoso-manganic 
oxide  —  the  form  in  which  it  is  customary  to  weigh  —  is  too 
uncertain.  It  had  been  supposed,  also,  that  the  presence  of 
ammoniacal  salts  (as  well  as  of  carbonic  acid)  causes  solution 
of  the  manganous  carbonate,  until  the  work  of  Guyard  (Hugo 
Tamm)  f  showed  that  when  the  precipitation  is  accomplished 
by  ammonium  carbonate,  even  hi  the  presence  of  ammonium 
chloride,  complete  separation  of  the  manganese  is  possible. 
No  data  are  given  by  Guyard  to  show  the  completeness  of  the 
separation  of  the  manganese  by  this  process ;  but  Fresenius  J 
examined  the  method  and  speaks  favorably  of  it.  In  this 
process  the  main  difficulty  of  the  older  method  of  estimation 
as  the  carbonate  —  viz.,  the  inclusion  of  the  alkaline  salt  —  is 
avoided.  We  know  now  how  to  avoid  the  difficulty  in  the 
way  of  weighing  as  the  oxide  by  converting  that  substance  to 
the  form  of  the  sulphate,  as  shown  in  a  former  paper.  § 

For  a  careful  study  of  the  separation  of  manganous  carbonate 
by  Guyard's  method  a  solution  of  pure  manganous  chloride 
was  prepared  and  standardized  as  the  anhydrous  sulphate  in 

*  From  Am.  Jour.  Sci.,  v,  382.  t  Chem.  News,  xxvi,  37. 

J  Zeitschr.  anal.  Chem.,  xi,  290. 

§  Am.  Jour.  Sci.,  v,  209.    This  volume,  p.  77. 


ON  THE  ESTIMATION  OF  MANGANESE. 


97 


the  manner  detailed  in  the  paper  to  which  reference  has  been 
made  above.  A  definite  volume  of  the  manganous  chloride 
was  carefully  drawn  into  a  platinum  dish  and  diluted  to  a 
volume  of  200  cm3.  To  the  solution  heated  to  100°  C.  ammo- 
nium chloride  (about  10  grm.)  was  added  and  ammonium  car- 
bonate in  excess.  The  solution  was  kept  warm  until  the 
precipitate  subsided,  and  then  was  filtered  off  on  asbestos  on 
a  perforated  crucible  under  pressure.  The  presence  of  ammo- 
nium chloride  is  necessary  to  insure  such  a  condition  of  the 
precipitate  that  it  will  not  run  through  the  felt. 

Inasmuch  as  the  precipitate  was  collected  under  conditions 
which  readily  permit  an  attempt  to  weigh  as  the  carbonate,  a 
trial  of  that  method  was  made  incidentally.  The  event  proved, 
as  Rose  *  has  stated  previously,  that  when  the  carbonate  is 
gently  heated,  evolution  of  carbon  dioxide  and  oxidation  of 
the  residue  begins  before  the  water  is  thoroughly  removed; 
for,  though  nearly  all  the  results  are  above  the  theory,  the 
solution  of  the  residue  in  hydrochloric  acid  indicated  plainly 
the  presence  of  a  small  amount  of  a  higher  oxide  of  manganese. 
In  the  following  table  are  found  the  results  of  a  series  of 
experiments  in  which  the  attempt  was  made  to  weigh  first 
as  carbonate  and  again  after  strong  ignition  —  well  within  the 
oxidizing  flame  of  a  powerful  burner  f — as  the  manganoso- 
manganic  oxide.  The  application  of  the  bromine  test  to  the 
hot  ammoniacal  filtrate  showed  that  in  every  one  of  these 
experiments,  the  precipitation  of  the  manganese  in  the  form 
of  the  carbonate  had  been  complete. 


MnC03 

Mn304 

MnClz. 

MH4C1. 

Found. 

Theory. 

Error. 

Found. 

Theory. 

Error. 

cms 

grm. 

grm. 

grm. 

grm. 

grm. 

grm. 

grm. 

50 

10 

0.2685 

0.2680 

0.0005+ 

0.1770 

0.1776 

0.0006- 

50 

10 

0.2704 

0.0024+ 

0.1788 

0.0012+ 

50 

10 

0.2710 

0.0030+ 

0.1770 

0.0006- 

50 

10 

0.2720 

0.0040+ 

0.1774 

0.0002- 

*  Annal  Phys.  Chem.,  Ixxxiv,  52. 

t  Am.  Jour.  Sci.,  v,  209.    This  volume,  p.  82. 

VOL.   II.  —  7 


98 


ON  THE  ESTIMATION  OF  MANGANESE 


As  shown  in  this  table,  weighing  as  the  carbonate  is  out  of 
the  question;  the  errors  of  the  process  when  the  residue  is 
ignited  in  the  manner  described  to  form  the  manganoso- 
manganic  oxide  are  much  smaller  though  rather  variable. 
The  estimation  of  manganese  as  the  anhydrous  sulphate  had 
given  in  the  work  to  which  reference  has  been  made  results 
agreeing  so  much  more  closely  than  could  be  obtained  by 
any  other  procedure,  that  the  attempt  was  made  to  estimate 
the  amount  of  manganese  precipitated  as  the  manganese 
carbonate  by  converting  it  first  to  the  oxide,  then  to  the 
sulphate.  A  given  weight  of  sulphate  was  precipitated  as 
the  manganous  carbonate,  after  the  employment  of  all  the 
precautions  mentioned  previously  in  this  paper,  and  then 
filtered  off  on  ashless  filter  paper.  After  washing  thoroughly 
with  hot  water,  the  filter  was  burned,  the  residue  ignited  for 
the  condition  of  the  manganoso-manganic  oxide  and  weighed 
as  such.  Then  the  oxide  was  converted  to  the  sulphate  by 
heating  with  three  or  four  drops  of  concentrated  sulphuric 
acid.  The  agreement  of  the  results  as  shown  in  the  following 
table  is  considerably  better. 


Exp. 

MH4C1. 

Mn304 

MnSO4 

Found. 

Theory. 

Error. 

Found. 

Theory. 

Error. 

Ill 

18 

grm. 

10 
10 
10 
10 

grm. 

0.2463 
0.1110 
0.1584 
0.1672 

grm. 
0.2478 
0.1121 
0.1581 
0.1699 

grim* 
0.0015- 
0.0011- 
0.0003+ 
0.0027- 

grm. 
0.4903 
0.2225 
0.3126 
0.3355 

grm. 
0.4905 
0.2219 
0.3128 
0.3364 

grm. 
0.0002- 
0.0006+ 
0.0002- 
0.0009- 

By  treatment  of  the  filtrates  of  (1),  (2),  and  (3)  with 
bromine  and  ammonia  at  boiling  temperature  no  manganese 
was  found.  In  the  filtrate  from  number  (4)  by  the  same 
treatment,  a  small  amount  of  manganese  dioxide  was  precipi- 
tated, which  when  heated  with  concentrated  sulphuric  acid 
gave  0.0006  grm.  of  manganous  sulphate;  and,  hence,  the 
error  in  that  determination  is  really  0.0003  grm.  on  the 


SEPARATED  AS   THE   CARBONATE.  99 

sulphate.     The  slightly  larger  deficiency  recorded  in  the  table 
was  probably  due  to  imperfect  filtering. 

It  seems  to  be  evident  that  Guyard's  method  of  separation 
of  manganese  as  the  manganous  carbonate,  when  handled  with 
precautions,  gives  complete  separation  of  that  element.  It 
must  be  recognized  clearly,  however,  that  the  precipitation 
should  be  made  in  the  presence  of  a  considerable  amount  of 
ammonium  chloride,  and  that  great  care  must  be  used  in  the 
filtering  and  washing  of  the  finely  divided  precipitate.  It  is 
altogether  preferable  to  weigh  in  the  condition  of  the  sulphate. 


XIV 

THE    ACTION    OF    CARBON    DIOXIDE   ON 
SOLUBLE   BORATES. 

BY  LOUIS  CLEVELAND  JONES* 

IN  a  process  for  the  separation  and  estimation  of  boric  acid 
devised  by  Morse  and  Burton,  f  the  liberation  of  carbonic, 
silicic,  and  boric  acids  from  a  mixture  of  inorganic  salts  is 
effected  by  the  action  of  sulphuric  acid,  using  tropseolin  0  0 
as  an  indicator  of  acidity.  In  the  solution  thus  prepared,  con- 
taining in  free  condition  only  carbonic,  silicic  and  boric  acids, 
the  silicic  acid  is  dehydrated  and  made  insoluble  by  anhydrous 
copper  sulphate.  The  boric  acid  is  then  extracted  with  abso- 
lute alcohol.  To  this  alcoholic  solution  of  boric  acid,  a  known 
amount  of  barium  hydroxide  solution  is  added  in  excess  over 
that  required  to  form  a  barium  metaborate,  BaB2O4.  Carbon 
dioxide  is  then  passed  into  the  solution  in  accordance  with  the 
hypothesis  that  the  excess  only  of  barium  is  acted  upon.  The 
aqueous  mixture  of  barium  metaborate  and  barium  carbonate 
is  evaporated  and  the  residue  is  heated  to  a  constant  weight 
over  a  triple  burner.  From  the  following  proportion  the  boric 
acid  present  may  be  calculated.  The  molecular  weight  of 
boric  acid  —  the  molecular  weight  of  carbon  dioxide  :  the 
molecular  weight  of  boric  acid  =  the  total  weight  found  —  the 
theoretical  weight  of  barium  as  carbonate  :  the  weight  of  boric 
acid  present.  It  is  obvious,  inasmuch  as  the  difference  be- 
tween the  calculated  weight  of  the  barium  as  carbonate  and 
the  actual  weight  of  the  residue  is  multiplied  nearly  three  times 
to  get  the  boric  oxide,  that  the  actual  error  of  the  process, 
whatever  it  may  be,  is  magnified  threefold  by  the  method  of 
computation. 

*  From  Am.  Jour.  Sci.,  v,  442.  t  Am.  Chem.  Jour.,  x,  154. 


ACTION  OF  CARBON  DIOXIDE  ON  BORATES.      101 


I  have  made  a  study  of  this  method  applied  to  pure  boric 
acid,  but  have  been  unable  to  obtain  results  similar  to  those  of 
Morse  and  Burton. 

For  this  investigation  boric  acid  of  standard  strength  was 
made  by  dissolving  in  a  given  amount  of  water  a  known 
weight  of  anhydrous  boric  oxide,  prepared  by  igniting  over  a 
blast  lamp  boric  acid  several  tunes  recrystallized  and  washed. 
A  solution  of  barium  hydroxide  was  filtered  free  from  carbon- 
ate and  then  standardized  by  precipitation  as  carbonate  and 
also  by  the  Phelps  method  with  iodine.*  To  a  measured 
amount  of  the  boric  acid  solution  an  excess  of  barium  hydrox- 
ide was  added,  carbon  dioxide  passed,  and  the  whole  evapo- 
rated and  by  successive  ignitions  brought  to  a  constant  weight. 
Below  are  tabulated  some  of  the  results : 


Exp. 

Ba(OH)8 

taken. 
Calculated 
as  BaCO,. 

B203 

taken. 

Weight  of  residue 
after  successive 
ignitions. 

B203 
found. 

Error. 

grin,. 

grin. 

grm. 

grm. 

grm. 

(1) 

0.9391 

0.2200 

(    1st  wt.,  0.9860 
}    2d    "    0.9786 

0.1263 
0.1063 

0.0937- 
0.1137- 

(    1st    "    0.9605 

0.0744 

0.0523- 

(2) 

0.9318 

0.1295 

1   2d    "    0.9558 

0.0646 

0.0649- 

(    3d    "    0.9510 

0.0517 

0.0778- 

1st   "    1.0357 

0.2972 

0.0780+ 

2d    "    1.0248 

0.2679 

0.0487+ 

(3) 

0.9253 

0.2192 

- 

8d    "    1.0149 
4th  "    1.0064 

0.2412 
0.2183 

0.0220+ 
0.0009- 

5th  "    0.9975 

0.1944 

0.0248- 

6th  "    0.9855 

0.1621 

0.0671- 

1st   "    0.8017 

0.1927 

0.1117+ 

2d    "    0.7777 

0.1281 

0.0471+ 

3d    "    0.7642 

0.0918 

0.0108+ 

4th  "    0.7582 

0.0757 

0.0053- 

(4) 

0.7301 

0.0810 

- 

5th  "    0.7517 
6th  "    0.7482 

0.0582 
0.0487 

0.0228- 
0.0323- 

7th  "    0.7447 

0.0393 

0.0417- 

8th  "    0.7427 

0.0339 

0.0471- 

9th  "    0.7422 

0.0326 

0.0484- 

10th  "    0.7407 

0.0285 

0.0525- 

Plainly  the  results  vary  with  the  degree  of  the  ignition. 
At  the  outset  the  residue  may  or  may  not  weigh  more  than 
the  theory  requires  for  the  known  amounts  of  barium  hydrox- 
*  Am.  Jour.  Sci.,  ii,  70.    Volume  I,  p.  369. 


102  ACTION  OP  CARBON  DIOXIDE 

ide  and  boric  acid  taken  upon  the  assumption  that  the  residue 
is  barium  metaborate  and  barium  carbonate.  This  is  obviously 
natural  if  the  carbon  dioxide  acts  upon  the  barium  borate  as 
well  as  upon  the  excess  of  barium  hydroxide  ;  for,  it  is  to  be 
expected  that  in  the  evaporation  more  or  less  of  the  free  boric 
acid  will  volatilize,  and  that  in  the  subsequent  ignition  the 
boric  acid  remaining  will  tend  to  recombine  more  or  less  com- 
pletely, replacing  carbon  dioxide.  If  the  boric  acid  present 
were  to  recombine  completely  with  the  barium  carbonate  to 
form  a  metaborate,  the  final  result  would  always  be  low  by 
just  the  amount  of  free  boric  acid  volatilized  in  the  process 
of  evaporation  and  ignition.  The  evidence  of  an  experiment, 
however,  in  which  0.25  grm.  of  previously  prepared  barium 
metaborate  was  fused  in  contact  with  0.5  grm.  of  barium  car- 
bonate, resulting  in  a  loss  of  0.0871  grm.,  goes  to  show  that 
the  metaborate  and  carbonate  of  barium  interact  still  further 
to  liberate  carbon  dioxide. 

These  results  were  so  surprising  in  the  light  of  the  experi- 
ence of  Morse  and  Burton  that  the  question  of  the  possibility 
of  breaking  up  by  carbon  dioxide  the  barium  metaborate 
already  formed  was  put  to  the  test  directly.  A  known  amount 
of  barium  hydroxide  was  taken  in  solution  and  to  it  added  an 
amount  of  boric  acid  very  little  in  excess  of  that  theoretically 
necessary  to  form  the  barium  metaborate.  The  solution  was 
evaporated  to  dryness  and  the  residue  ignited.  The  weight 
obtained  proved  to  be  0.0008  grm.  less  than  the  sum  of  the 
barium  and  boric  oxides  taken,  doubtless  because  the  slight 
excess  of  boric  acid  was  somewhat  volatile  in  the  evaporation. 
The  mass,  presumably  barium  metaborate,  was  now  dissolved 
as  completely  as  possible  in  hot  water,  carbon  dioxide  was 
passed  through  the  solution,  the  whole  was  evaporated,  and 
the  residue  ignited  and  weighed.  The  increase  in  the  weight 
showed  that  carbon  dioxide  had  been  absorbed,  while  a  corre- 
sponding amount  of  boric  acid  had  not  volatilized. 


grm.  Krm-  Prm-  grm. 

(5)      0.8377      0.2990      0.9491      0.9771 


OAT  SOLUBLE  BORATES. 


103 


After  passing  in  carbon  dioxide  and  igniting  the  increase  in 
weight  was  0.0280  grm.,  representing  the  gas  absorbed  less 
the  boric  acid  volatilized. 

It  was  plain  that  barium  metaborate  is  decomposed  in  solu- 
tion by  carbon  dioxide.  The  possibility  remained,  however, 
that  the  action  of  carbon  dioxide  might  be  so  regulated  as  to 
leave  the  metaborate  practically  unattacked.  In  experiment 
(6),  therefore,  carbon  dioxide  was  passed  above  the  stirred 
solution  until  no  further  precipitate  formed  upon  the  surface, 
the  barium  present  being  in  excess  of  that  required  to  form  a 
metaborate. 


Ba(OH)2  taken.        ,,  n  .  .          Weight  of  residue 
Calculated  as  BaC03.  B*°s  taken'        after  ignition. 


(6) 


grm. 

0.7094 


grin. 

0.2070 


grm. 

0.7710 


BjO8  found, 
grm. 

0.1658 


Error. 

grm. 
0.0412- 


The  variation  of  this  result  from  the  theory  shows  that 
under  these  conditions  the  metaborate  is  not  unaffected  by 
carbon  dioxide,  the  loss  being  due,  of  course,  to  the  escape  of 
boric  acid. 

An  attempt  was  now  made  therefore  to  gauge  the  amount 
of  the  carbon  dioxide  introduced  by  means  of  an  indicator. 
In  experiment  (7)  phenolphthalein  was  added  to  the  solution 
of  boric  acid  containing  an  excess  of  barium  hydroxide  and 
the  current  of  gas  was  stopped  when  the  color  of  the  indi- 
cator disappeared. 


Exp. 

» 

Calculated 

B20S 
taken. 

Weight  of  residue 
after  ignition. 

found. 

Error. 

as  BaCO3. 

grm. 

grm. 

grm. 

grill. 

grm. 

1st  wt.,  0.6820 

0.2011 

0.0438+ 

2d    "     0.6783 

0.1911 

0.0338+ 

3d    "     0.6720 

0.1742 

0.0169+ 

(7) 

0.6073 

0.1573 

. 

4th  "     0.6700 

0.1688 

0.0115+ 

5th  "    0.6655 

0.1567 

0.0006- 

6th  "     0.6630 

0.1499 

0.0074- 

7th  "    0.6609 

0.1443 

0.0130- 

This  result  is  manifestly  an  improvement  over  those  obtained 
without  the  careful  restriction  of  the  supply  of  carbon  dioxide. 


104  ACTION  OF  CARBON  DIOXIDE 

A  similar  experiment,  differing  only  in  the  single  point  that 
the  carbon  dioxide  was  made  to  act  upon  a  boiling  solution, 
resulted  in  like  manner. 

In  the  light  of  these  observations  it  is  plain  that  a  sufficiently 
prolonged  action  of  carbon  dioxide  should  result  in  the 
displacement  of  all  the  boric  acid  if  that  acid  can  be  removed 
from  the  field  of  action  as  fast  as  it  is  liberated.  Experiments 
were  made  which  clearly  demonstrate  the  truth  of  this 
hypothesis.  A  small  side-necked  flask  was  charged  with  a 
solution  of  boric  acid  (0.1143  grm.)  and  barium  hydroxide 
(0.3227  grm.)  in  proportion  to  form  the  metaborate.  The 
mass  was  brought  nearly  to  dryness  by  distillation  and  methyl 
alcohol  (15  cm3)  added.  Through  this  flask,  in  which  the 
alcohol  was  kept  boiling  by  a  Bunsen  burner,  was  passed  the 
vapor  of  methyl  alcohol,  while  carbon  dioxide,  purified  by  a 
neutral  solution  of  silver  nitrate,  bubbled  continually  through 
the  entire  system.  The  methyl  alcohol  vapor  coming  from  the 
side-neck  flask  was  kept  lighted  by  contact  with  the  flame  of 
a  Bunsen  burner  and  the  distillation  continued  for  two  hours 
until  the  flame  showed  not  the  slightest  tinge  of  green.  The 
residue  in  the  flask  originally  containing  barium  metaborate 
was  brought  to  dryness  and  tested  for  boric  acid.  Only  a 
trace  was  found,  and  this  was  thought  to  be  due  to  inclusion 
by  the  insoluble  barium  carbonate.  In  a  similar  experiment 
in  which  borax  was  used  instead  of  barium  borate,  no  trace  of 
boric  acid  was  found  in  the  residue  of  sodium  carbonate,  either 
by  turmeric  or  the  flame  test,  while  the  distillation  was 
continued  only  one-half  the  time  of  the  preceding  experiment. 
This  result  is  quite  in  harmony  with  the  views  of  P. 
George  vie*  to  the  effect  that  the  large  absorption  of  carbon 
dioxide  in  solutions  of  borax  indicates  that  the  boric  acid  is 
displaced  from  its  union  with  the  base. 

Obviously,  the  division  of  the  base  between  boric  acid  and 
carbonic  acid  falls  under  the  principle  of  mass  action,  and  if 
the  boric  acid  is  taken  in  sufficient  excess  over  the  barium 
hydroxide,  the  action  of  the  carbonic  acid  should  be  inappreci- 

*  Jour.  prak.  Chem.  xxxviii,  118. 


ON  SOLUBLE  BORATES.  105 

able.  This  idea  is  sustained  by  an  experiment  in  which  about 
1  grm.  of  boric  acid  was  dissolved  with  0.15  grm.  of  barium 
hydroxide.  No  precipitate  could  be  obtained  by  passing 
carbon  dioxide  and  boiling.  In  fact,  in  the  French  process  of 
borax  manufacture  just  this  action  of  an  excess  of  boric  acid 
upon  a  boiling  solution  of  sodium  carbonate  is  used.  On  the 
other  hand  to  prevent  the  formation  of  a  carbonate  this  excess 
of  boric  acid  must  be  considerable  if  the  action  of  carbon 
dioxide  is  prolonged.  Thus  in  an  experiment  in  which  a 
current  of  carbon  dioxide  was  passed  into  a  solution  containing 
0.1219  grm.  of  boric  anhydride  and  one-half  the  amount  of 
barium  needed  to  form  a  metaborate,  the  solution  (60  cm3) 
deposited  on  boiling  90  per  cent  of  the  barium  present  in  the 
form  of  the  carbonate. 

In  view  of  the  facts  which  I  have  described,  it  is  difficult  to 
see  under  what  conditions  Morse  and  Burton  prevented  the 
excessive  action  of  carbon  dioxide  and  obtained  hi  their 
analytical  method  the  excellent  results  which  they  record. 


XV 

FURTHER   SEPARATIONS   OF  ALUMINUM  BY 
HYDROCHLORIC   ACID. 

BY  FRAME  STUART  HAVENS.* 

IN  former  papersf  from  this  laboratory  methods  have  been 
described  for  the  separation  of  aluminum  from  ferric  iron,  and 
from  beryllium,  based  on  the  insolubility  of  the  hydrous 
aluminum  chloride  in  a  mixture  of  equal  parts  of  aqueous 
hydrochloric  acid  and  ether  saturated  with  hydrochloric  acid 
gas,  the  ferric  and  beryllium  chlorides  being  exceedingly 
soluble  in  this  mixture. 

It  was  the  purpose  of  the  investigations  herein  described  to 
discover  how  far  this  process  could  be  extended,  with  certain 
modifications,  to  cover  the  separation  of  aluminum  from  such 
other  metals  as  might  occur  with  it,  either  in  artificially 
prepared  alloys  or  in  naturally-occurring  compounds. 

The  aluminum  used  in  all  the  following  experiments  was  in 
the  form  of  a  solution  of  the  chloride.  This  chloride  was 
purified,  as  previously  described, J  from  iron  by  precipitation 
with  hydrochloric  acid,  and  from  the  alkalies  by  precipitation 
as  the  hydroxide  and  continued  washing  with  water  until  the 
washings  gave  no  test  with  silver  nitrate.  The  hydroxide 
thus  obtained  was  dissolved  in  hot  hydrochloric  acid  to  get  it 
into  the  form  of  the  chloride.  The  chloride  solution  was 
standardized  by  precipitating  weighed  portions  with  ammonia 
and  weighing  as  the  oxide. 

*  From  the  Am.  Jour.  Sci.,  vi,  45. 

t  Gooch  and  Havens,  Am.  Jour.  Sci.,  vol.  ii,  p.  416.     This  volume,  p.  20 ; 
Havens,  Am.  Jour.  Sci.,  vol.,  iv,  p.  111.    This  volume,  p.  47. 
J  Loc.  cit. 


FURTHER  SEPARATIONS  OF  ALUMINUM,  ETC.       107 

Separation  of  Aluminum  from  Zinc. 

A  solution  of  pure  zinc  chloride  made  by  dissolving  metallic 
zinc,  free  from  impurities,  in  hydrochloric  acid,  is  not 
precipitated  when  treated  with  an  equal  volume  of  ether 
and  saturated  with  hydrochloric  acid  gas. 

To  prepare  a  definite  zinc  salt  free  from  traces  of  the  alkalies 
which  would  be  precipitated  with  the  aluminum  by  strong 
hydrochloric  acid,  pure  metallic  zinc  was  dissolved  in  hydro- 
chloric acid,  the  dilute  solution  precipitated  with  ammonium 
carbonate,  and  the  resulting  carbonate  ignited  to  a  constant 
weight  as  zinc  oxide.  This  oxide  dissolved  in  hydrochloric 
acid  gave  a  pure  chloride. 

The  aluminum  in  all  these  experiments  was  determined  in 
the  following  manner: — Portions  of  the  prepared  solution  of 
aluminum  chloride  were  weighed  in  a  small  beaker,  weighed 
portions  of  zinc  oxide  added  and  sufficient  aqueous  hydrochloric 
acid  to  dissolve  it.  The  beaker  was  then  cooled  by  immersion 
in  an  inverted  bell-jar  supplied  with  running  water  by  means 
of  inlet  and  outlet  tubes,  and  a  current  of  gaseous  hydrochloric 
acid  (generated  by  the  gradual  addition  of  sulphuric  acid  to  a 
mixture  of  hydrochloric  acid  and  salt)  passed  through  the 
solution  in  the  beaker  to  complete  saturation.  Ether  was 
added  in  volume  equal  to  that  of  the  original  solution,  and  the 
whole  again  saturated  with  hydrochloric  acid  gas.  The 
crystalline  chloride  precipitated  was  caught  on  asbestos  in  a 
filter  crucible,  washed  with  a  previously  prepared  solution  of 
equal  parts  of  ether  and  hydrochloric  acid,  saturated  with 
hydrochloric  acid  gas,  dried  for  half  an  hour  at  150°-180°  C., 
covered  with  a  layer  of  pure  mercury  oxide,  heated  gently 
over  a  low  flame  under  a  ventilating  hood,  ignited  over  the 
blast,  and  weighed  as  the  oxide.  The  results  show  that 
aluminum  can  be  determined  with  reasonable  accuracy  in  the 
presence  of  a  pure  zinc  salt. 

The  zinc  can  be  determined,  after  the  evaporation  of  the 
strong  acid  filtrate,  by  any  of  the  well  known  methods.  It 
was  found,  however,  that  after  thorough  conversion  to  the 


108 


FURTHER  SEPARATIONS  OF  ALUMINUM 


nitrate  by  repeated  evaporation  with  nitric  acid  the  salt  could 
be  ignited  directly  to  the  oxide  with  satisfactory  results.  This 
is  shown  clearly  in  Table  I,  (3)  to  (5).  In  experiments 
(3)  and  (4),  zinc  oxide  was  dissolved  in  nitric  acid  and  the 
nitrate  ignited  again  to  the  oxide.  In  experiment  (5),  the  zinc 
oxide  was  first  dissolved  in  hydrochloric  acid  and  the  chloride 
thus  obtained  was  converted  to  the  nitrate  by  evaporating  the 
solution  (5  cm3)  with  nitric  acid  (2  cm3),  treating  the  residue 
with  nitric  acid  (2  cm3)  and  evaporating  to  dryness. 

TABLE  I. 


Bxp. 

A12O3  taken 
as  the 
chloride. 

found. 

Error. 

ZnO 
taken. 

ZnO 
found. 

Error. 

Error 
corrected. 

Final, 
vol. 

gnu. 

grin. 

grm. 

grm. 

grm. 

grm. 

grm. 

cm8 

(1) 

0.0562 

0.0562 

0.0000 

0.1110 

. 

> 

. 

(2) 

0.0580 

0.0577 

0.0003- 

0.1034 

_ 

t 

t 

(3) 

f 

. 

. 

0.1019 

0.1016 

0.0003- 

t 

§ 

(4) 

,  . 

.  . 

.  . 

0.1010 

0.1007 

0.0003- 

.  . 

, 

(5) 

.  . 

.  . 

0.1100 

0.1095 

0.0005- 

(6) 

0.0572 

0.0572 

0.0000 

0.1014 

0.1027 

0.0013+ 

0.0007- 

12 

(7) 

0.0563 

0.0550 

0.0013- 

0.1026 

0.1038 

0.0012+ 

0.0008- 

16 

8) 

0.0577 

0.0576 

0.0001- 

0.1000 

0.1014 

0.0014+ 

0.0006- 

16 

(9) 

0.0559 

0.0558 

0.0001- 

0.1020 

01035 

0.0015+ 

0.0005- 

16 

(10) 

0.0563 

0.0556 

0.0007- 

0.2024 

0.2046 

0.0022+ 

0.0002+ 

20 

(11) 

0.1111 

0.1107 

0.0004- 

0.2092 

0.2116 

0.0024+ 

0.0004+ 

20 

In  Table  I,  (6)  to  (11),  are  given  the  results  of  experiments 
in  which  both  the  aluminum  and  zinc  were  determined,  —  the 
former,  as  described,  by  precipitating  as  the  hydrous  chloride 
and  weighing  as  the  oxide,  and  the  latter  by  carefully  evapo- 
rating the  strongly  acid  filtrate  (best  with  a  small  current  of 
air  playing  on  the  surface  of  the  liquid  to  avoid  spattering  due 
to  the  too  violent  evolution  of  the  ether  and  gaseous  acid)  and 
finally  converting  the  chloride  through  the  nitrate  into  the 
oxide.  It  is,  of  course,  absolutely  necessary  that  the  treatment 
with  nitric  acid  shall  be  thorough,  so  that  no  zinc  chloride  may 
remain  to  volatilize  when  the  residue  is  ignited.  On  account 
of  the  danger  to  platinum  from  the  aqua  regia  generated  by 
the  action  of  nitric  acid  on  zinc  chloride,  the  evaporations  of 
the  filtrates  from  the  aluminum  chloride  and  the  treatment 


BY  HYDROCHLORIC  ACID. 


109 


with  nitric  acid  were  carried  on  in  porcelain  and  the  residual 
nitrate  was  transferred  to  a  small  crucible  for  ignition.  In 
this  process  the  porcelain  was  evidently  attacked  somewhat, 
so  that  the  residual  nitrate  was  slightly  contaminated  with 
material  from  the  large  porcelain  dish.  This  fact  accounts 
for  the  high  results  given  in  the  first  column  of  errors. 
However,  on  introducing  a  correction  (0.0020)  found  by 
carrying  through  the  process  in  blank  with  the  quantities  of 
reagents  employed  in  the  regular  process,  the  results  on  zinc, 
slightly  deficient,  agree  closely  with  those  obtained  in  (3)-(5), 
Table  I,  where  the  zinc  nitrate  was  converted  directly  to 
the  oxide  without  the  previous  evaporation  in  porcelain  of 
a  large  volume  of  strongly  acid  liquid.  The  errors  thus 
corrected  stand  in  another  column  of  the  table. 

These  results  show  clearly  that  aluminum  and  zinc  may  be 
separated  from  one  another  by  the  action  of  hydrochloric  acid 
gas  in  aqueous  ethereal  solution  with  a  reasonable  degree  of 
accuracy. 

Separation  of  Aluminum  from  Copper,  Mercury,  and  Bismuth. 

The  separation  of  aluminum  from  copper,  mercury,  and 
bismuth  does  not  differ  materially  from  the  separation  of 

TABLE  II. 


Exp. 

A12O8  taken 
as  the 
chloride. 

found. 

Error. 

CuO 

taken. 

CuO 
found. 

Error. 

HgCl, 
taken. 

Bi203 

taken. 

gj 

(3) 

18 

(6) 

gnu. 

0.0576 
0.0561 
0.0570 
0.0548 
0.0565 
0.0576 

gnn. 
0.0571 
0.0557 
0.0574 
0.0557 
0.0571 
0.0577 

gnn. 
0.0005- 
0.0004- 
0.0004+ 
0.0009+ 
0.0006+ 
0.0001+ 

gnn. 
0.0500 
0.0400 

gnu. 

gnn. 

gTIH. 

0.1000 
0.1000 

grin. 

o.iobo 

0.2000 

J3 

(9) 
10) 

11) 
12) 
13) 

0.0558 
0.0538 
0.0566 
0.0577 

0.0545 
0.0536 
0.0562 
0.0575 

0.0013- 
0.0002- 
0.0004- 
0.0002- 

0.0437 
0.0359 
0.0345 
0.0319 
0.0343 
0.0337 
0.0651 

0.0432 
0.0359 
0.0340 
0.0324 
0.0356 
0.0349 
0.0644 

0.0005- 
0.0000 
0.0005- 
0.0005+ 
0.0013+ 
0.0012+ 
0.0007- 

110      FURTHER  SEPARATION  OF  ALUMINUM,  ETC. 

aluminum  and  zinc.  Aluminum  chloride  is  precipitated 
quantitatively  in  the  presence  of  pure  salts  of  these  elements 
as  shown  in  experiments  of  Table  II. 

In  determining  the  copper  in  the  acid  filtrate  it  was  found 
advantageous  to  weigh  as  the  oxide,  but  to  arrive  at  that 
condition  through  the  sulphate  rather  than  through  the  nitrate 
(which  was  the  transition  salt  in  the  case  of  zinc),  as  this 
process  can  be  carried  on  safely  in  platinum. 

In  Table  II,  (10)-(13),  are  given  results  of  experiments  in 
which  the  aluminum  was  determined  as  previously  described 
by  precipitation  as  the  hydrous  chloride  and  conversion  to  the 
oxide.  The  acid  nitrate  was  evaporated  in  platinum  and  the 
copper  determined  by  treating  the  residue  with  a  few  drops  of 
strong  sulphuric  acid,  heating  gently  to  drive  off  the  excess  of 
sulphuric  acid,  and  then  igniting  the  sulphate  to  the  oxide  at 
a  red  heat.  That  the  copper  sulphate  is  converted  to  the 
oxide  by  ignition  at  a  red  heat  over  a  Bunsen  burner  is  shown 
in  experiments  (7)  to  (9)  of  Table  II. 


XVI 

THE  IODOMETRIC  DETERMINATION  OF 
MOLYBDENUM. 

BY  F.  A.  GOOCH  AND  JOHN  T.  NORTON  JR.* 

A  PROCESS  for  the  iodometric  determination  of  molybdic  acid, 
which  consists  in  treating  a  soluble  molybdate  in  a  Bunsen 
distillation-apparatus  with  potassium  iodide  and  hydrochloric 
acid,  has  been  advocated  by  Friedheim  and  Euler.f  Accord- 
ing to  this  process  the  molybdate,  containing  from  0.2  grm. 
to  0.3  grm.  of  molybdenum  trioxide,  is  treated  with  from 
0.5  grm.  to  0.75  grm.  of  potassium  iodide  and  enough  hydro- 
chloric acid,  of  sp.  gr.  1.12,  to  fill  two-thirds  of  the  flask  of 
the  apparatus.  The  liquid  is  warmed  until  heavy  vapors  of 
iodine  fill  the  flask  and  then  boiled  until  iodine  vapor  is  no 
longer  visible  and  the  color  of  the  liquid  residue  is  a  clear 
green.  The  iodine  liberated  is  collected  in  the  distillate 
and  titrated  with  sodium  thiosulphate,  every  atom  of  iodine 
found  indicating  presumably  the  reduction  of  a  molecule  of 
molybdic  acid  to  the  condition  of  the  pentoxide  Mo2O5. 

It  was  pointed  out  in  a  former  article  from  this  laboratory,  t 
that  greater  precaution  than  was  taken  by  Friedheim  and 
Euler  is  necessary  in  order  that  the  reduction  may  proceed 
according  to  theory,  and  that  the  iodine  collected  may  serve 
as  a  reliable  measure  of  the  molybdic  acid.  It  was  found 
that  the  green  color  of  the  liquid  comes  gradually  and  that 
it  may  develop  distinctly  before  the  molybdic  acid  is  fully 
reduced.  It  was  found,  also,  that  since  even  a  trace  of 
oxygen  liberates  iodine  from  the  hot  mixture  of  potassium 

*  From  Am.  Jour.  Sci.,  vi,  168. 

t  Ber.  Dtsch.  chem.  Ges.,  xxviii,  2066. 

t  Gooch  and  Fairbanks,  Am.  Jour.  Sci.,  ii,  156.    Volume  I,  p.  375. 


112  THE  IODOMETRIC  DETERMINATION 

iodide  and  hydrochloric  acid  of  the  strength  employed,  it  is 
not  sufficient  to  rely  upon  the  volatilization  of  iodine  to 
expel  the  air  originally  in  the  apparatus,  but  that  it  is 
essential  to  conduct  the  distillation  in  an  atmosphere  devoid 
of  oxygen.  The  suggestion  was  made  therefore  that  the 
operation  should  be  carried  on  in  a  current  of  carbon  dioxide 
and  that  the  mixture,  constituted  definitely,  should  be  boiled 
between  stated  limits  of  concentration  which  were  determined 
by  experiment.  It  was  found  that  when  amounts  of  a  soluble 
molybdate  containing  less  than  0.3  grm.  of  molybdenum 
trioxide  are  treated  with  potassium  iodide,  not  exceeding  the 
theoretical  proportion  by  more  than  0.5  grm.,  and  40  cm3  of 
a  mixture  of  the  strongest  hydrochloric  acid  and  water  in 
equal  parts,  the  reduction  proceeds  with  a  fair  degree  of 
regularity  and  is  practically  complete  when  the  volume  has 
diminished  to  25  cm3.  If  the  operation  is  properly  conducted 
in  an  atmosphere  of  carbon  dioxide,  it  was  shown  that  the 
iodine  in  the  distillate  may  be  trusted  to  indicate  the  molybdic 
acid  within  reasonable  limits  of  accuracy.  It  appeared,  how- 
ever, that  too  great  an  excess  of  potassium  iodide  tends  to 
induce  excessive  reduction,  and  that  the  same  tendency 
shows  when  the  liquid  is  concentrated  to  too  low  a  limit. 

To  this  criticism  Friedheim  took  exception  *  and  contrasted, 
to  their  disadvantage,  our  results  by  the  modified  method 
with  those  of  Friedheim  and  Euler  by  the  original  method. 
It  became  necessary,  therefore,  to  point  out  f  the  fact  that 
of  the  results  published  by  Friedheim  and  Euler,  upon  which 
reliance  was  placed  to  prove  the  reliability  of  their  method, 
five  out  of  seven  in  one  series  and  one  out  of  five  in  another 
series  had  been  calculated  incorrectly  from  data  given. 
Another  series  of  six  determinations  was,  however,  apparently 
faultless  in  this  respect.  More  recently  f  Euler  has  explained 
that  the  errors  were  not  really  arithmetical.  Two  of  them 
may  be  presumed,  inferentially,  to  be  due  to  careless  copying 

*  Ber.  Dtsch.  chem.  Ges.,  xxix,  2981. 
t  Gooch,  Am.  Jour.  Sci.,  iii,  237. 
J  Zeit.  anorg.  Chem.,  xv,  454. 


OF  MOLYBDENUM.  113 

or  proof-reading;  and  four,  we  are  told  by  Euler,  were 
introduced  into  the  series  by  mistake,  and  actually  represent 
(as  Prof.  Friedheim  kindly  informs  him)  the  analysis  of  a 
sample  of  ammonium  molybdate  of  undetermined  constitution : 
that  is  to  say,  the  figures  now  given  by  Euler  represent  the 
original  percentages  of  molybdenum  trioxide  which  had  been 
changed  by  some  unconscious  process  from 

80.62  per  cent  to  81.85  per  cent. 
80.71       «         "  81.69       « 

80.63  "         "  81.67       " 
80.78       "         "  81.78       « 

Curiously  enough,  Euler's  corrected  figures,  as  given  here, 
are  still  affected  by  trifling  arithmetical  errors  of  from  one  to 
four  units  in  the  second  decimal  place.  The  agreement  of 
these  results  among  themselves  is  no  proof  of  the  correctness 
of  the  process  of  analysis.  The  great  variation  between  the 
average  percentage  of  molybdenum  trioxide  in  ammonium 
molybdate  as  found  by  Euler  in  a  molybdate  of  known  con- 
stitution and  the  percentage  of  the  trioxide  as  found  by 
Friedheim  (if  we  understand  Euler  aright)  may  be  due  con- 
ceivably to  either  or  both  of  two  causes,  viz. :  the  change  of 
material  analyzed,  and  the  change  of  operator  or  conduct  of 
the  operation.  We  shall  show  in  the  following  account  of 
our  work  that  the  exact  control  of  the  conditions  of  treatment, 
along  the  lines  laid  down  formerly,  is  actually  essential  to  the 
reduction  of  molybdic  acid  according  to  the  theory  of  the 
process. 

Our  experiments  were  made  with  ammonium  molybdate 
twice  recrystallized  from  the  presumably  pure  salt.  The  con- 
stitution of  the  preparation  was  determined  by  careful  ignition 
per  se,  and,  for  greater  security,  with  sodium  tungstate  free 
from  carbonate.  It  contained  81.83  per  cent  of  molybdenum 
trioxide. 

The  potassium  iodide  which  we  used  was  prepared  by  act- 
ing with  re-sublimed  iodine  upon  iron  wire,  and  precipitating 
by  potassium  carbonate  —  the  proportions  of  iodine  and  iron 

VOL.    II.  — 8 


114 


THE  IODOMETRIC  DETERMINATION 


having  been  adjusted  to  secure  the  formation  of  the  hydrous 
magnetic  oxide  of  iron.  The  filtrate  from  the  iron  hydroxide 
gave  on  evaporation  and  crystallization  potassium  iodide  which 
was  free  from  iodate. 

The  hydrochloric  acid  was  taken  of  sp.  gr.  1.12,  because 
this  is  the  strength  used  by  Friedheim  and  Euler. 

The  sodium  thiosulphate  employed  was  taken  in  nearly 
decinormal  solution,  and  was  standardized  by  running  it  into 
an  approximately  decinormal  solution  of  iodine  which  had 
been  determined  by  comparison  with  decinormal  arsenious 
acid  made  from  carefully  re-sublimed  arsenious  oxide.  We 
chose  this  method  of  standardizing  —  the  introduction  of  the 
thiosulphate  into  the  iodine  —  rather  than  the  reverse  opera- 
tion, in  order  that  the  conditions  of  the  actual  analysis  might 
be  followed  in  the  standardization. 

The  distillation  apparatus  was  constructed  with  sealed  or 
ground  joints  of  glass  wherever  contact  with  iodine  was  a 

possibility.  It  was  made  by 
sealing  together  a  separating 
funnel  A,  a  100  cm3  Voit 
flask  B,  a  Drexel  wash-bottle 
C,  and  a  bulbed  trap  g,  as 
shown  in  the  figure.  Upon 
the  side  of  the  distillation- 
flask  B  was  pasted  a  gradu- 
ated scale  by  means  of  which 
the  volume  of  the  liquid 
within  the  flask  might  be 
known  at  any  time.  Carbon 
dioxide,  generated  in  a  Kipp 
apparatus  by  the  action  of 

dilute  hydrochloric  acid  (carrying  in  solution  cuprous  chloride 
to  take  up  free  oxygen)  upon  marble  previously  boiled  in 
water,  was  passed  through  the  apparatus  before  and  during 
the  operation,  so  that  it  was  possible  to  interrupt  the  process 
of  boiling  at  any  point  of  concentration,  to  remove  the  receiver 
by  easy  manipulation,  to  replace  the  receiver,  and  to  continue 


FIG.  20. 


OF  MOLYBDENUM.  115 

the  distillation  without  danger  of  admitting  air  to  the  distilla- 
tion flask. 

In  experiments  to  be  described  —  (1)  to  (5)  of  the  table  — 
the  proportions  of  potassium  iodide  and  molybdic  acid,  and 
the  strength  of  the  hydrochloric  acid  recommended  by  Fried- 
heim  and  Euler  were  retained.  The  essential  change  of 
condition  is  the  removal  of  atmospheric  air  from  the  distillation 
flask  before  the  acid  is  admitted  to  contact  with  the  other 
reagents.  Potassium  iodide  (3  grm.)  and  water  (200  cm3) 
were  put  into  the  receiver  C,  and  a  little  of  this  solution  was 
allowed  to  flow  into  the  trap  g.  Ammonium  molybdate  care- 
fully weighed  (0.3  grm.)  and  potassium  iodide  (0.5  grm.  to 
0.75  grm.)  were  introduced  into  the  distillation  flask  B,  the 
apparatus  was  connected  as  shown  in  the  figure,  and  carbon 
dioxide  was  passed  freely  through  the  whole  apparatus  for 
some  minutes.  The  stop-cock  d,  between  the  bulb  of  the 
funnel  A  and  the  flask  B,  was  closed,  and  hydrochloric  acid 
(40  cm3,  sp.  gr.  1.12)  was  poured  into  the  funnel ;  the  air 
above  the  liquid  in  the  funnel  was  displaced  by  carbon  dioxide 
through  the  space  between  the  neck  of  the  funnel  and  the 
loosely  adjusted  stopper  carrying  the  inlet  tube ;  the  connec- 
tion between  the  funnel  and  inlet  tube  was  tightened,  the 
stop-cock  opened,  and  the  acid,  under  the  pressure  of  carbon 
dioxide,  was  permitted  to  flow  into  the  flask.  In  this  way 
the  acid,  iodide,  and  molybdate  were  made  to  interact  with 
little  danger  of  the  presence  of  oxygen.  The  flask  was  heated 
by  the  Bunsen  burner,  and  the  iodine  evolved,  passing  over 
quietly  in  the  slow  current  of  carbon  dioxide,  collected  in  the 
receiver.  The  liquid  was  boiled  until  fumes  of  iodine  were 
no  longer  visible  above  the  liquid  in  the  flask  and  connecting 
tubes  backed  by  a  ground  of  white,  and  then  a  full  minute 
more.  At  this  stage,  the  green  color  of  the  liquid  having 
developed  fully,  the  apparatus  was  permitted  to  cool,  the 
current  of  carbon  dioxide  was  increased,  the  cap  of  the  receiver 
was  loosened  at  /,  the  contents  of  the  trap  were  washed  back 
into  the  receiver,  the  rest  of  the  apparatus  was  lifted  bodily 
from  the  receiver,  the  liquid  adhering  to  the  inlet  tube  was 


116  THE  IODOMETRIC  DETERMINATION 

washed  off  into  the  receiver,  and  the  end  of  the  tube  was 
dipped  immediately  into  a  solution  of  potassium  iodide.  The 
constant  flow  of  carbon  dioxide  prevented  reflux  of  air  during 
the  transfer,  and  as  soon  as  the  end  of  the  tube  had  been 
submerged  in  the  solution  of  potassium  iodide  (which  was 
employed  not  only  as  a  water-seal,  but  to  catch  any  iodine 
still  carried  in  the  gas),  it  was  possible  to  reduce  the  rapidity 
of  the  current. 

After  titrating  the  iodine  in  the  distillate  the  receiver  was 
again  placed  in  the  train  and  the  process  of  distillation  was 
resumed  under  the  former  conditions  and  continued  until  the 
volume  of  the  liquid,  as  indicated  upon  the  scale,  had  dimin- 
ished to  25  cm3,  when  the  distillation  was  interrupted.  The 
apparatus  was  manipulated  as  before  to  prevent  access  of  air, 
and  the  iodine  evolved  in  the  second  treatment  determined.  A 
third  period  of  distillation  served  to  show  the  iodine  liberated 
during  the  concentration  of  the  liquid  from  25  cm3  to  10  cm3. 

During  the  first  period  of  distillation  the  liquid  assumed  the 
clear  green  color,  which  changed  but  slightly  until  the  begin- 
ning of  the  third  period,  when  the  tint  verged  upon  olive,  and 
at  the  end  of  the  operation  the  color  of  the  liquid  was  an 
olive  brown  which  grew  browner  on  cooling.  The  addition 
of  considerable  hydrochloric  acid  to  the  residual  liquid  re- 
stored the  clear  green  color,  while  water  changed  the  olive 
brown  to  reddish  yellow,  the  tint  varying  with  the  dilution. 
The  results  of  these  experiments  are  recorded  in  (1)  to  (5)  of 
the  accompanying  table.  In  division  A  are  given  the  weights 
of  molybdenum  trioxide  corresponding  to  the  amounts  of 
iodine  found  in  the  three  stages  of  distillation ;  in  division  B, 
the  molybdenum  trioxide  corresponding  to  the  iodine  evolved 
from  the  beginning  of  the  process  to  the  end  of  each  stage. 

The  mean  error  of  the  indications  taken  during  the  period 
of  distillation  advocated  by  Friedheim  and  Euler  is  0.0045 
grm.  — ;  *  that  of  the  period  of  concentration  from  40  cm3  to 

*  Even  this  figure  does  not  disclose  the  full  error,  which  is  partly  counter- 
balanced, as  will  appear  later,  by  the  effect  of  oxygen  dissolved  in  the  acid 
used  in  the  process. 


OF  MOLYBDENUM. 


11T 


A. 

Exp. 

HCl 

Sp&i12 

Klin 

retort. 

MoOs  taken 
as  ammonium 
molybdate. 

Mo03  corresponding  to  iodine  found. 

First  stage 
40cm3  to 
32cm3. 
Green  color. 

Second  stage 
32cm3  to 
25  cm3. 

Third  stage 
25cm3  to 
10cm3. 

(1) 
(2) 
(3) 
(4) 
(5) 

cm8 
40 
40 
40 
40 
40 

grm. 

0.5 
0.5 
0.5 
0.75 
0.75 

grm. 
0.2455 
0.2455 
0.2455 
0.2455 
0.2455 

grm. 
0.2399 
0.2402 
0.2414 
0.2404 
0.2431 

grm. 
0.0076 
0.0053 
0.0040 
0.0061 
0.0037 

grm. 

0.0004 
0.0013 
0.0004 
0.0004 
0.0004 

(6) 
(7) 

40 
40 

1 

2 

0.2455 
0.2455 

0.2404 

0.0085 

0.0019 

B. 

Ezp. 

MoO3  corre- 
sponding to 
iodine  found 
during  period 
of  Friedheim 
and  Euler. 
(1st  stage.) 

Error. 

MoO3  corre- 
sponding to 
iodine  found  in 
concentrating 
from  40cm3 
to  25cm3. 

Error. 

Mo03  corre- 
sponding to 
iodine  found  in 
concentrating 
from  40cm3 
to  10  cm3. 

Error. 

(1) 
(2) 
(3) 
(4) 
(5) 

grm. 

0.2399 
0.2402 
0.2414 
0.2404 
0.2431 

grin. 
0.0056- 
0.0053- 
0.0041- 
0.0051- 
0.0024- 

grin. 

0.2475 
0.2455 
0.2454 
0.2465 
0.2468 

gnu. 
0.0020+ 
0.0000 
0.0001- 
0.0010+ 
0.0013+ 

grm. 
0.2479 
0.2468 
0.2458 
0.2469 
0.2472 

grm. 
0.0024+ 
0.0013+ 
0.0003+ 
0.0014+ 
0.0017+ 

(6) 
(7) 

0.2404 

0.0051- 

0.2489 

0.0034+ 

0.2508 
(  0.2495 
\  0.2529* 

0.0053+ 
0.0040+ 
0.0074+ 

25  cm3  is  0.0008  grm.  +  ;  and  that  of  the  full  period  of  distilla- 
tion is  0.0014  +.  It  is  plain  beyond  a  peradventure  that  in 
the  process  as  conducted  by  Friedheim  and  Euler,  except- 
ing the  protection  against  atmospheric  action  the  theoretical 
reduction  of  the  molybdic  acid  does  not  take  place.  The  best 
results  are  obtained  when  the  distillation  is  prolonged  until 
the  original  volume  of  40  cm3  has  been  diminished  to  25  cm8. 
Concentration  beyond  the  limit  of  25  cm3  tends  to  develop  a 
tendency  toward  over-reduction,  especially  when  the  amount 
of  potassium  iodide  is  increased  beyond  about  0.5  grm.  in 

*  On  repeating  distillation  with  a  fresh  charge  of  acid. 


118  THE  IODOMETRIC  DETERMINATION 

excess  of  that  theoretically  required.  This  is  shown  in  exper- 
iments (6)  and  (7),  conducted  otherwise  similarly  to  those 
described  above,  in  which  the  amount  of  potassium  iodide  was 
increased  to  1  grm.  and  2  grms.  The  error  after  distilling 
from  40  cm3  to  10  cm3,  the  lowest  limit  of  the  preceding  ex- 
periments, was  0.0053  grm.  +  and  0.0040  grm.  -f,  and  the 
latter  error  was  increased  to  0.0074  grm.  -f  on  repeating 
the  distillation  with  a  fresh  portion  (30  cm3)  of  the  acid.  It 
is  interesting  to  note  incidentally  that  in  the  experiment  in 
which  the  largest  amount  of  iodide  (2  grms.)  was  used  the 
solution  did  not  take  the  green  color  at  any  stage  of  the  dis- 
tillation, probably  because  the  large  excess  of  iodide  held  the 
free  iodine  and  so  masked  the  color  until  the  degree  of  con- 
centration was  reached  at  which  the  olive-brown  color  dis- 
places the  green. 

The  possibility  of  the  interaction  of  atmospheric  oxygen 
and  gaseous  hydriodic  in  the  analytical  process,  even  to  the 
extent  of  producing  errors  of  from  one  to  three  per  cent 
reckoned  as  molybdenum  trioxide,  was  recognized  by  Fried- 
heim  and  Euter ;  and  it  was  to  obviate  this  difficulty  that  the 
recommendation  was  made  by  them  to  warm  very  gradually 
the  distillation  flask  filled  two-thirds  with  the  mixture  of 
iodide,  molybdate,  and  acid,  and  to  raise  the  liquid  to  actual 
boiling  only  when  the  space  above  the  liquid  in  the  retort  and 
in  the  connecting  tube  is  filled  as  completely  as  possible  with 
iodine  vapor,  while  the  liquid  in  the  receiver  begins  to  rise  in 
the  tube. 

The  action  of  atmospheric  oxygen  upon  the  solution  of  hydri- 
odic acid  must,  however,  be  also  taken  into  account.  It  is  a 
familiar  fact  that  when  a  considerable  excess  of  strong  hydro- 
chloric acid  is  allowed  to  act  in  contact  with  air  upon  potassium 
iodide  (free  from  iodate)  dissolved  in  a  little  water,  the  mix- 
ture is  colored  by  free  iodine.  The  amount  of  iodine  liberated 
by  atmospheric  action  is  insignificant  when  the  acid  is  very  di- 
lute, but  is  considerable  when  the  acid  is  strong,  and  increases 
with  tune  and  rise  in  temperature,  as  shown  in  the  experiments 
recorded  in  the  accompanying  table. 


OF  MOLYBDENUM. 


119 


Per- 

KI 

taken. 

Volume. 

centage 
of  HC1 
in 
aqueous 

Time 
in 
minutes. 

Temperature. 
Centigrade. 

MoOs 
equivalent 
to  iodine 
found. 

Remarks. 

acid. 

gnu. 

cm8 

gnu. 

1 

66 

2 

1 

23° 

None. 

1 

66 

2 

10 

23° 

0.0001 

1 

66 

24* 

10 

23° 

0.0017  ^ 

Diluted  to  500  cm3 

1 

1 

66 
66 

24* 
24* 

4 
10 

(  From  23°  to 
J   the  boiling 
(    point. 

0.0067  [ 
0.0121  ) 

before     titrating 
with  Na2S208. 

Even  the  precaution  to  conduct  the  operation  in  an  atmos- 
phere of  carbon  dioxide  does  not  eliminate  all  chance  of  error 
of  this  sort  unless  the  liquid  of  the  mixture  —  the  hydro- 
chloric acid  —  is  free  from  air.  The  experiments  of  the  fol- 
lowing statement,  which  were  conducted  in  the  apparatus  and 
manner  previously  described,  show  this  point  clearly.  Thus, 


Per- 

TUff\f\^ 

KI 

taken. 

VoL 

centage 
of  HC1  in 
aqueous 
acid. 

Concentration 
by 
boiling. 

Jjfl.OUo 

equivalent 
to  iodine 
found. 

Remarks. 

grm. 

cm8 

cm8. 

grm. 

(  40-30 

0.0013 

Iodine  determined 

1 

40 

24 

} 

in  distillate. 

(30-20 

0.0002 

1  grm.  of  KI  added 

to  retort  at  the 

beginning  of  the 

2d  stage. 

1 

40 

20 

40-25 

0.0005 

The    acid    taken, 

sp.  gr.  1.1,  was 

freshly      boiled 

and    introduced 

at  once  upon  KI 
in  retort  in  CO2. 

40  cm3  of  unboiled  acid,  sp.  gr.  1.12,  introduced  enough  air 
into  the  apparatus  to  cause  an  error  of  0.0013  grm.  reckoned 
in  terms  of  molybdenum  trioxide,  while  the  iodine  set  free  by 
the  action  of  the  residual  acid  of  this  experiment  upon  another 
gram  of  potassium  iodide  introduced  without  admission  of  air 
corresponded  to  only  0.0002  grm.  in  terms  of  molybdenum 

*  This  corresponds  nearly  to  sp.  gr.  1.12. 


120     IODOMETRIC  DETERMINATION  OF  MOLYBDENUM. 

trioxide.  The  use  of  acid  of  sp.  gr.  1.1,  freshly  boiled  in  the 
air,  obviously  reduces  the  error  due  to  the  unboiled  acid,  but 
even  in  this  case  the  effect  of  included  oxygen  was  not  wholly 
obviated. 

It  is  obvious  that  the  procedure  recommended  by  Friedheim 
and  Euler  can  by  no  possibility  eliminate  the  effect  of  atmos- 
pheric action  upon  the  mixture  of  acid  and  iodide.  The 
extent  of  such  action  must  depend  upon  such  conditions  as  the 
size  of  the  apparatus,  the  time  of  exposure,  the  body  of  air 
above  and  dissolved  in  the  liquid,  and  the  rate  of  displacement 
of  the  air.  How  great  the  error  due  to  atmospheric  action 
actually  was  in  the  process  as  conducted  by  Friedheim  and 
Euler  we,  of  course,  have  no  means  of  knowing.  It  is  to  be 
hoped,  however,  that  it  was  sufficiently  great  to  counterbalance 
that  other  inevitable  error  (of  about  five  milligrams)  which 
exists  by  reason  of  the  incompleteness  with  which  molybdic 
acid  is  reduced  under  the  conditions  which  these  investigators 
prescribe ;  for,  the  value  of  Euler's  work  upon  the  vanadio- 
molybdates  rests  upon  the  chance  that  these  two  very 
considerable  and  indisputable  tendencies  to  error  may  have 
neutralized  one  another. 

It  has  been  shown  clearly  that  our  former  criticism  of  the 
procedure  of  Friedheim  and  Euler  is  justified  in  every  par- 
ticular. We  have  no  change  to  make  in  the  recommendation 
made  therein  as  to  necessary  modifications. 

If  the  conditions  seem  difficult,  there  is  an  alternative  in 
the  method  proposed  in  the  former  article,*  according  to 
which  the  molybdate  is  reduced  by  the  acid  and  iodide  in  an 
Erlenmeyer  beaker  (trapped  loosely  by  means  of  a  short 
bulbed  tube  hung  in  the  neck)  and  the  molybdenum  pentoxide, 
freed  from  iodine  by  boiling,  is  reoxidized  by  standard  iodine 
in  alkaline  solution. 

*  Am.  Jour.  Sci.,  ii,  156.    Volume  I,  p.  375. 


XVII 

ON  THE   DETERMINATION  OF  MANGANESE 
AS  THE  PYROPHOSPHATE. 

BY  F.  A.   GOOCH  AND  MAKTHA  AUSTIN.* 

FOR  the  estimation  of  manganese  in  a  gravimetric  way  when 
accuracy  is  a  consideration,  recourse  is  usually  taken  to  the 
excellent  method  of  Wolcott  Gibbs.  f  This  method  consists  in 
the  precipitation  of  a  manganous  salt  by  an  alkaline  phosphate, 
the  conversion  of  the  tribasic  phosphate  into  the  ammonium 
manganese  phosphate,  and  the  weighing  of  the  product  of 
ignition  as  the  pyrophosphate. 

By  Gibbs'  original  method  the  orthophosphate  of  manganese 
was  precipitated  by  hydrogen  disodium  phosphate  in  large 
excess  above  the  quantity  required  to  cause  the  precipitation. 
The  flocky  white  precipitate  was  dissolved  either  in  sulphuric 
or  hydrochloric  acid,  and  precipitated  again  at  the  boiling 
temperature  by  ammonia  in  excess.  This  semi-gelatinous 
precipitate  on  boiling  or  long  standing  even  in  the  cold  becomes 
crystalline,  the  crystals  forming  beautiful  talcose  scales  which 
have  a  pearly  luster  and  a  pale  rose  color.  The  precipitate 
was  filtered  off,  washed  with  hot  water,  dried  and  ignited.  The 
results  obtained  by  Gibbs'  students  for  the  pyrophosphate 
accord  closely  with  the  theory. 

Fresenius  J  showed  subsequently  that  ammonium  manganese 
phosphate  dissolves  in  cold  water,  in  hot  water,  and  in  an 
aqueous  solution  of  ammonium  chloride  [1  :  70]  to  the  extent 
of  1  part  in  32,000,  1  part  in  20,000,  and  1  part  in  18,000, 
respectively.  It  is  clear,  however,  that  the  solubility  of  this 

*  From  Am.  Jour.  Sci.,  vi,  233. 
t  Am.  Jour.  Sci.,  xliv,  216. 
J  Zeitschr.  anal.  Chem.,  vi,  415. 


122  DETERMINATION  OF  MANGANESE 

precipitate  is  not  indicated  necessarily  by  the  proportions 
given  so  long  as  an  excess  of  the  precipitant  is  present  during 
the  washing,  though  Fresenius  did  find  in  the  filtrate  traces 
of  manganese  which  to  his  mind  were  sufficient  to  account 
for  losses  indicated  by  his  test  analyses,  viz.,  one  to  three 
milligrams  of  oxide,  or  from  two  to  six  milligrams  of 
phosphate. 

Another  mode  of  manipulation  has  been  advocated  by  Blair  * 
in  order  that  the  precipitate  may  be  obtained  more  easily  in 
crystalline  condition.  According  to  this  method  dilute  ammonia 
is  added  drop  by  drop  to  the  hot  acid  solution  until  the  precipi- 
tate begins  to  form,  the  boiling  and  stirring  are  continued 
until  the  small  amount  of  flocky  precipitate  is  converted 
completely  to  crystalline  condition,  and  the  process  of  adding 
ammonia  drop  by  drop  is  repeated  until  the  manganese  is  all 
down  in  crystalline  condition.  The  dilute  ammonia  is  added 
in  excess  and  the  liquid  filtered  after  cooling  in  ice  water. 

In  discussing  these  methods  of  precipitation,  McKenna  f 
points  out  that  both  give  good  and  accordant  results,  and  that 
the  process  may  be  carried  on  in  glass  as  well  as  in  platinum, 
if  the  time  of  crystallization  is  made  short  enough. 

When  a  manganous  salt  is  precipitated  in  the  cold  by  an 
excess  of  an  alkaline  phosphate,  it  falls,  as  Heintz  f  has  shown, 
in  the  form  of  the  trimanganous  phosphate  of  the  formula 
Mn3P2O8.  This  same  phosphate  constitutes,  as  we  have  found, 
the  greater  part  of  the  precipitate  which  forms  when  a  man- 
ganous salt  reacts  in  the  cold  in  the  presence  of  ammonium 
chloride  with  microcosmic  salt  and  ammonia  in  slight  excess. 
Boiling  or  even  subsequent  standing  may,  as  is  well  known, 
effect  a  more  or  less  complete  conversion  of  the  manganese 
phosphate  to  the  ammonium  manganese  phosphate.  Thus,  in 
one  experiment  hi  which  an  amount  of  manganous  chloride 
enough  to  produce  0.2214  gram  of  the  pyrophosphate  was 
precipitated  in  the  cold  by  5  cm3  of  a  saturated  solution  of 
microcosmic  salt,  with  the  subsequent  addition  of  ammonia 

*  The  Chemical  Analysis  of  Iron,  106.         t  Jour.  Anal.  Chem.,  v,  141. 
i  Ann.  Phys.,  cl,  449. 


AS   THE  PYROPHOSPHATE.  123 

in  excess,  in  a  volume  of  200  cm3  containing  also  5  grams  of 
ammonium  chloride,  the  residue  after  ignition  weighed  0.1904 
gram.  Presuming  this  residue  to  consist  entirely  of  the 
pyrophosphate  and  the  trimanganous  orthophosphate,  the  pro- 
portion of  the  former  to  the  latter  calculated  from  the  relation 
of  symbols,  and  the  weights  taken  and  found,  is  nearly  one  to 
six.  That  is  to  say,  about  six-sevenths  of  the  precipitate  fell 
in  this  experiment  in  the  form  of  the  tribasic  orthophosphate. 
In  another  experiment  made  exactly  similarly,  excepting 
that  the  .liquid  was  heated  to  boiling,  the  proportion  of  the 
manganese  pyrophosphate  to  the  trimanganous  orthophosphate 
in  the  only  partially  crystallized  precipitate  proved  to  be  two 
to  one.  That  is,  in  this  case,  two-thirds  of  the  precipitate  was 
in  the  form  of  the  pyrophosphate.  In  the  former  of  the 
experiments  a  small  amount  of  manganese  was  found  in  the 
filtrate,  but  not  enough  to  change  materially  the  ratio  recorded. 
The  slight  solubility  appears  to  be  connected  with  the  incomplete 
conversion  of  the  trimanganous  phosphate  to  the  ammonium 
manganese  phosphate,  for  as  will  appear  later,  the  manganese 
found  in  the  nitrate,  when  the  conversion  is  known  to  be 
nearly  complete,  is  inappreciable  unless  extraordinary  amounts 
of  the  ammonium  salt  are  present.  The  success  of  the  analytical 
process  under  discussion  turns,  therefore,  upon  the  change  of 
the  trimanganous  phosphate  Mn3P2O8  to  the  ammonium  man- 
ganese phosphate  NH4MnPO4.  In  the  work  to  be  described 
the  attempt  was  made  to  learn  the  conditions  under  which  this 
conversion  may  be  best  and  most  completely  accomplished. 

The  conversion  of  a  molecule  of  trimanganous  phosphate 
to  the  ammonium  manganese  phosphate  might  be  due,  con- 
ceivably, either  to  the  action  of  free  ammonia  or  to  the  action 
of  a  salt  of  ammonium.  The  action  of  ammonia  could  only 
take  place  at  the  expense  of  a  partial  loss  of  manganese  from 
the  phosphate  and  its  appearance  as  a  hydroxide,  two-thirds 
of  the  manganese  going  into  two  molecules  of  the  ammonium 
manganese  phosphate.  In  the  presence  of  ammonium  salts 
it  is  possible  that  the  manganese  oxide  thus  replaced  might 
enter  into  union  with  the  acid  radical  of  the  ammonium  salt 


124  DETERMINATION  OF  MANGANESE 

setting  free  ammonia ;  but  if  the  ammonium  salt  present  were 
the  phosphate,  or  if  an  alkaline  phosphate  were  present  with 
other  suitable  ammonium  salts,  it  is  conceivable  that  the 
replaced  manganese  might  appear  as  a  constituent  of  a  third 
molecule  of  ammonium  manganese  phosphate.  In  any  event, 
it  would  be  the  ammonium  salt  and  not  the  free  ammonia 
which  would  determine  the  formation  of  the  third  molecule 
of  the  ammonium  manganese  phosphate.  Plainly,  too,  the 
ammonium  salt  by  itself,  if  it  were  a  phosphate,  or  if  a  soluble 
phosphate  were  also  present,  might  accomplish  the  conversion 
without  the  intermediate  action  of  free  ammonia.  Unless, 
therefore,  free  ammonia  favors  the  insolubility  of  the  ammo- 
nium manganese  phosphate,  its  presence  would  be  unnecessary 
and  might  even  be  an  actual  disadvantage  if  the  hydroxide 
naturally  formed  by  its  action  upon  the  manganese  phosphate 
were  to  fail  to  reunite  fully  with  a  phosphoric  acid  radical. 
It  is  plain,  too,  that  the  action  of  free  ammonia  might  not 
stop  with  the  replacement  of  one  out  of  the  three  of  the 
manganese  atoms  present  in  the  molecule,  but  might  even 
proceed  under  favorable  conditions  to  the  formation  of  phos- 
phate richer  in  ammonium  and  to  the  separation  of  more 
manganese  from  its  union  with  the  acid  radical.  As  a  matter 
of  fact  Munroe  *  has  shown  that  the  prolonged  action  of  hot 
ammonia  upon  the  precipitate  produced  by  the  interaction  of 
a  manganous  salt  and  an  alkaline  phosphate  does  actually 
produce  a  hydroxide  which  blackens  as  it  takes  oxygen  from 
the  air.  Our  attention  has  been  given,  therefore,  more 
especially  to  a  study  of  the  conditions  of  action  under  which 
a  salt  of  ammonium  —  the  chloride  —  may  bring  about  the 
conversion  of  the  precipitate  first  thrown  down  by  an  alkaline 
phosphate  to  the  form  of  the  ammonium  manganese  phosphate. 
Experiments  were  made  upon  solutions  of  pure  manganous 
chloride  prepared  and  standardized  by  means  of  the  sulphate 
method,  as  described  in  a  former  paper,  f  to  show  the  effect 
of  varying  amounts  of  ammonium  chloride  on  the  condition 

*  Amer.  Chemist,  1877. 

t  Am.  Jour.  ScL,  v,  209.    This  volume,  p.  77. 


AS   THE  PYROPHOSPHATE. 


125 


of  the  precipitate  and  upon  the  solubility  of  the  precipitate 
when  once  formed.  The  ammonium  chloride  for  this  work 
was  prepared  pure  by  boiling  the  chemically  pure  salt  of 
commerce  with  a  faint  excess  of  ammonium  hydrate  and 
filtering  —  to  free  it  from  traces  of  iron,  silica  and  alumina. 
In  the  first  series  of  experiments  dilute  ammonia  was  added 
slowly  to  the  hot  faintly  acidulated  solution  containing  the 
manganous  chloride  and  more  than  enough,  theoretically,  of 
a  saturated  solution  of  microcosmic  salt  to  precipitate  the 
manganese  present.  The  liquid  was  heated  and  stirred  until 
the  flocky  mass  was  changed  to  a  crystalline  condition.  The 
addition  of  ammonia  drop  by  drop,  with  constant  stirring  and 
heating,  was  continued  until  the  manganese  was  all  precipi- 
tated in  crystalline  form.  A  slight  excess  of  ammonia  was 
added  and  the  liquid  with  the  precipitate  was  allowed  to 
stand  for  a  half  hour,  cooling  gradually  or  chilled  in  ice 
water.  The  precipitate  was  filtered  off  on  asbestos  under 
pressure,  washed  carefully  in  water  made  faintly  ammoniacal, 
dried  and  ignited.  The  filtrates  were  tested  for  manganese 
by  treatment  with  bromine  and  heating.  The  results  of  these 
experiments  are  given  in  the  following  table. 

TABLE  I. 


Mn2P2O7  equivalent 
to  MnCl,. 

Error 
in  terms  of 
Mn2P207. 

Error  in 
terms  of 
Manganese. 

Saturated 
solution  of 
HNH4NaPO4. 
4H20. 

Total 
volume. 

Manganese 
in  nitrate. 

Taken. 

Found. 

grm. 

grm. 

gnn. 

grm. 

cm» 

cm3 

0.4033 

0.3769 

0.0264- 

0.0102- 

5 

60* 

None. 

0.4033 

0.3728 

0.0305- 

0.0118- 

5 

60* 

None. 

0.3770 

0.3530 

0.0240- 

0.0090- 

5 

60 

None. 

0.3770 

0.3620 

0.0150- 

0.0058- 

5 

60 

None. 

0.4033 

0.3751 

0.0282- 

0.0109- 

10 

60 

None. 

0.4033 

0.3774 

0.0259- 

0.0100- 

10 

60 

None. 

0.4033 

0.3871 

0.0162- 

0.0062- 

5 

200 

None. 

0.3226 

03066 

0.0160- 

0.0062- 

5 

200 

None. 

In  this  method  of  precipitation  of  the  manganese  in  a  pure 
solution  of  a  manganous  salt  the  results  are  all  wrong.     The 


*  Chilled  in  ice-water. 


126  DETERMINATION  OF  MANGANESE 

proportion  of  the  trimanganous  phosphate  to  the  pyrophos- 
phate  in  the  residue,  calculated  from  the  symbols  and  the 
weights  taken  and  found,  is  in  the  average  two  to  five. 
That  is.  to  say,  five-sevenths  of  the  trimanganous  phosphate 
has  been  converted  to  the  form  of  the  ammonium  manganous 
phosphate. 

The  precipitate  obtained  in  this  manner  is  white  and  granu- 
lar but  not  silky,  and  after  ignition  it  shows  the  same  dead 
white  color,  and  is  powdery.  Evidently  the  regulation  of  the 
volume  in  which  the  precipitation  is  made  is  not  essential, 
and  the  chilling  of  the  liquid  is  of  no  importance  in  changing 
the  manganese  to  the  ammonium  manganese  salt  under  the 
given  conditions.  It  is  plain,  moreover,  that  the  assumption 
of  a  crystalline  condition  cannot  serve  as  an  indication  that  the 
composition  of  the  salt  is  ideal.  It  is  to  be  noted,  however, 
that  the  conditions  obtaining  here  are  essentially  different 
from  those  hi  common  practice ;  for,  ordinarily,  when  man- 
ganese is  to  be  determined  ammonium  salts  are  abundantly 
present  as  the  result  of  previous  steps  in  analysis. 

In  the  experiments  of  the  next  series  the  conditions  are 
varied  simply  in  this  respect,  that  ammonium  salts  are  in- 
troduced before  the  precipitation.  The  precipitate  was  less 
granular  and  more  silky.  After  ignition  the  mass  was  white 
with  a  faint  rose  color.  In  the  experiments  of  section  A  of 
the  table  the  precipitate  first  thrown  down  was  redissolved, 
reprecipitated  and  filtered  after  cooling;  in  those  of  section 
B,  the  precipitate  was  filtered  after  cooling  without  re-solution 
and  without  reprecipitation ;  and  hi  those  of  section  C,  the 
first  precipitate  was  filtered  at  once  while  the  solution  was 
still  hot.  The  length  of  digestion  before  filtering  and  the 
indications  of  manganese  in  the  filtrate  are  recorded  in  the 
table. 

It  was  observed  in  these  experiments  that  when  the  amount 
of  ammonium  chloride  is  present  in  considerable  quantity  a 
fine  crystalline  condition  is  got  much  more  readily  than  when 
the  amount  of  that  salt  is  small:  with  maximum  amounts  of 
ammonium  chloride  the  change  from  the  flocky  to  the  crystal- 


AS   THE  PYROPHOSPHATE. 
TABLE  IL 


127 


Mn2P,OT  equiva- 
lent to  the  MnClj. 

Error  in 
terms  of 
Mn,P,O7. 

Error  in 
terms  of 
Mangan- 
ese. 

Saturated 
solution  of 
HNaNH4P04. 
4H,0. 

NH4C1. 

Total 
volume. 

Time  of 
stand- 

a. 

Mangan- 
ese in  the 
filtrate. 

Taken. 

Found. 

A. 

grm. 

grm. 

grm. 

grm. 

cm3 

grm. 

cm3 

hrs. 

0.1542 

0.1520 

0.0022- 

0.0008- 

5 

5 

200 

15 

None. 

0.1542 

0.1540 

0.0002- 

0.0000 

5 

10 

200 

15 

None. 

0.1542 

0.1536 

0.0006- 

0.0002— 

5 

10 

100 

5 

None. 

0.1542 

0.1535 

0.0007- 

0.0002- 

5 

20 

200 

21 

None. 

0.3770 

0.3712 

0.0058- 

0.0022— 

5 

20 

200 

i 

None. 

0.3770 

0.3724 

0.0046- 

0.0018- 

5 

20 

200 

1 

None. 

0.3084 

0.3069 

0.0015- 

0.0006- 

5 

40 

200 

1 

None. 

0.3084 

0.3060 

0.0024- 

0.0009- 

5 

40 

200 

1 

None. 

0.3084 

0.3059 

0.0025- 

0.0009- 

5 

40 

200 

15 

Trace. 

0.3084 

0.3057 

0.0027- 

0.0010- 

5 

60 

200 

15 

None. 

B. 

0.1542 

0.1521 

0.0021- 

0.0008- 

5 

10 

100 

40 

None. 

0.1542 

0.1512 

0.0030- 

0.0010- 

5 

10 

200 

40 

None. 

0.1542 

0.1532 

0.0010- 

0.0003- 

6 

20 

200 

15 

None. 

0.1542 

0.1531 

0.0011- 

0.0004- 

5 

20 

100 

15 

None. 

0.3770 

0.3720 

0.0050- 

0.0019- 

5 

20 

200 

i 

None. 

0.3770 

0.3745 

0.0035- 

0.0014- 

5 

20 

200 

£ 

None. 

C. 

0.1542 

0.1519 

0.0023- 

0.0009- 

5 

16 

200 

None. 

0.1542 

0.1530 

0.0012- 

0.0004- 

5 

20 

200 

None. 

0.1542 

0.1525 

0.0017- 

0.0007- 

5 

30 

200 

None. 

0.3084 

0.3020 

0.0064- 

0.0025- 

5 

10 

200 

None. 

0.3084 

0.3053 

0.0031- 

0.0012- 

5 

20 

200 

None. 

0.3084 

0.3033 

0.0051- 

0.0020- 

5 

20 

200 

None. 

0.3084 

0.3039 

0.0045- 

0.0017- 

5 

60 

200 

Trace. 

line  condition  is  almost  immediate ;  even  in  the  cold  the 
change  takes  place  to  a  marked  extent  in  a  few  seconds.  No 
manganese  was  found  in  the  nitrate  by  boiling  with  bromine 
and  ammonia  —  a  test  which  is  capable  of  indicating  0.0001 
grm.  of  manganous  sulphate  in  500  cm3  of  water  containing 
60  grm.  of  ammonium  chloride  —  until  the  ammonium  chlo- 
ride amounted  to  20  per  cent  of  the  mass,  or  to  40  grm. 
in  200  cm3  of  the  liquid,  and  even  then  but  once  in  three  trials : 
even  when  the  proportion  was  30  per  cent  —  60  grm.  in  200 
cm3  —  the  solvent  action  of  the  ammonium  chloride  upon  the 


128 


DETERMINATION  OF  MANGANESE 


manganese  salt  was  trifling.  The  pyrophosphate  residues 
obtained  in  these  experiments,  as  well  as  in  all  those  recorded 
in  this  paper,  were  dissolved  in  nitric  acid  and  tested  for 
contamination  by  a  chloride ;  in  no  single  case  did  silver  nitrate 
produce  more  than  an  inappreciable  opalescence  in  the  solution. 
It  is  plain,  therefore,  that  the  variations  of  the  results  from 
theory  are  occasioned  by  variation  in  the  degree  of  conversion 
of  the  trimanganese  phosphate  to  the  ammonium  manganese 
phosphate,  and  that,  while  the  ammonium  chloride  shows  no 
appreciable  solvent  action  on  the  precipitate  in  the  presence  of 
the  precipitant,  its  effect  in  the  process  of  conversion  is  plainly 
evident.  For  the  smaller  amounts  of  the  manganese  salts 
(equivalent  to  0.1542  grm.  of  the  pyrophosphate)  the  effect 
of  the  ammonium  chloride  reaches  a  maximum  when  that  salt 
amounts  to  10  per  cent  of  the  solution ;  for  twice  that  amount 
of  manganese  salt,  the  best  results  were  obtained  by  doubling 
the  amounts  of  ammonium  chloride.  Either  line  of  treatment 
yields  under  the  most  favorable  conditions,  results  which  are 
passably  good,  but  the  advantage  inclines  slightly  to  the  first 
method  in  which  the  first  precipitate  was  dissolved  and 
reprecipitated  while  the  liquid  was  cooled  before  filtering. 

TABLE  IH. 


Mn2P2O7  equivalent 

Error  in 
terms 
of  MN2P2O7. 

Error 
in  terms 
of 
Manganese. 

Saturatecl 
solution  of 
HNaNH4 
P04  .  4H20. 

NH4C1. 

Total 
volume. 

Manganese 
in  the 
nitrate. 

Taken. 

Found. 

grin. 

grm. 

grin. 

grin. 

cm» 

grm. 

cm3 

0.2214 

0.2202 

0.0012- 

0.0005- 

5 

20 

200 

None. 

0.2214 

0.2202 

0.0012- 

0.0005- 

5 

20 

200 

None. 

0.2214 

0.2191 

0.0023- 

0.0009- 

5 

20 

200 

None. 

0.2214 

0.2191 

0.0023- 

0.0009- 

5 

20 

300 

None. 

0.2214 

0.2191 

0.0023- 

0.0009- 

5 

20 

300 

None. 

0.2214 

0.2185 

0.0029- 

0.0011- 

10 

20 

200 

None. 

0.2214 

0.2186 

0.0028- 

0.0010- 

20 

20 

300 

None. 

0.2214 

0.2192 

0.0022- 

0.0009- 

20 

20 

300 

None. 

In  Table  III  are  recorded  results  obtained  by  precipitating 
the  cold  acid  solution  of  the  manganese  salt  and  the  microcos- 
mic  salt  with  a  strong  excess  of  ammonia.  The  mixture  was 


AS   THE  PYROPHOSPHATE. 


129 


heated  to  boiling  for  from  five  to  ten  minutes  and  filtered  hot. 
In  this  series  of  determinations  the  amount  of  ammonium 
chloride  present  was  constant  while  the  volume  of  the  liquid 
present  was  varied  and  the  amounts  of  the  microcosmic  salt. 

These  results  are  possibly  a  trifle  less  satisfactory  than  those 
obtained  for  the  smaller  amounts  of  manganese  by  the  method 
of  Table  II,  it  may  be  because  the  prolonged  boiling  tends  to 
form  a  trifling  amount  of  free  oxide  ;  but  the  fact  is  disclosed 
that  an  increase  of  the  microcosmic  salt  is  without  influence 
and  that  a  variation  in  volume  from  200  cm8  to  300  cm8  is  the 
occasion  of  little  change  in  the  indications  of  the  process. 

In  another  series  of  experiments  the  solution  of  manganous 
chloride  was  added  drop  by  drop  to  the  mixture  of  microcosmic 
salt  and  ammonium  chloride  made  alkaline  with  ammonia. 
The  precipitate  which  fell  in  the  cold  was  crystallized  by 
boiling  the  mixture  a  few  minutes.  The  results  are  given 
below: 

TABLE  IV. 


^oVctVaIent 

Error 

Saturated 

Error. 

in  terms 
of 
manganese. 

solution 
of 
HNaNH4P04. 

NH4C1. 

Total 
volume. 

in  the 
filtrate. 

Taken. 

Found. 

grm. 

grm. 

grm. 

grm. 

cms 

grm. 

cm.3 

0.1542 

0.1521 

0.0021- 

0.0008- 

5 

5 

200 

None. 

0.2214 

0.2203 

0.0011- 

0.0004- 

5 

10 

275 

None. 

0.2214 

0.2192 

0.0022- 

0.0009- 

5 

15 

275 

None. 

0.2214 

0.2197 

0.0017- 

0.0007- 

5 

20 

275 

None. 

0.2214 

0.2223 

0.0009+ 

0.0003+ 

5 

20 

200 

None. 

0.1542 

0.1528 

0.0014- 

0.0005- 

5 

30 

275 

None. 

The  experience  of  this  series  of  experiments  demonstrated 
again  that  the  ease  with  which  the  flocky  precipitate  is 
converted  to  the  crystalline  ammonium  manganese  phosphate 
is  proportioned  to  the  ammonium  chloride  present,  and  the 
mean  error  of  the  results  for  the  phosphate  when  the  ammonium 
chloride  reached  20  grams  (0.0007  grm.)  is  considerably  less 
than  the  mean  error  (0.0018  grm.)  when  the  amount  of  the 
ammonium  salt  was  less  than  20  grms. 

Experiments  were  also  made  according  to  the  modifications 

VOL.    II. — 9 


130  DETERMINATION  OF  MANGANESE 

suggested  by  Monroe,*  viz.,  the  boiling  of  the  manganous  salt 
with  an  excess  of  microcosmic  salt  until  the  precipitate  becomes 
crystalline  and  just  neutralizing  with  dilute  ammonia  ;  but  we 
have  been  unable  to  find  the  conditions  of  this  treatment  by 
which  uniform  results  may  be  obtained  in  even  moderate 
agreement  with  the  theory. 

We  have  tried  also  the  effect  of  substituting  ammonium 
nitrate  for  ammonium  chloride  in  the  conversion  process ;  but, 
so  far  as  our  experience  goes,  the  nitrate  is  not  so  effective 
weight  for  weight  in  producing  the  change  of  the  trimanganous 
phosphate  to  the  ammonium  manganese  phosphates,  while 
the  solubility  of  the  product  in  the  solution  of  the  ammonium 
nitrate  becomes  appreciable  more  rapidly  with  the  increase  of 
the  amount  present  than  is  the  case  when  the  ammonium  salt 
is  the  chloride. 

In  the  light  of  the  experiments  described  it  would  seem  to 
be  reasonable  to  expect  the  best  results  from  the  phosphate 
method  for  determining  manganese  when  the  conditions  are 
so  arranged  that  precipitation  may  take  place  in  the  cold 
solution  in  the  presence  of  but  little  free  ammonia,  and  of 
enough  ammonium  chloride  to  bring  about  the  rapid  conversion 
of  the  precipitate  to  the  crystalline  condition.  Under  such 
circumstances  it  should  be  possible  to  secure  the  conversion  of 
the  phosphate  to  the  ideal  constitution  as  completely  as  possible 
without  danger  of  subsequent  decomposition  by  the  prolonged 
action  of  the  hot  free  ammonia.  In  carrying  out  this  idea, 
the  solution  of  manganese  chloride  was  treated  as  before  with 
microcosmic  salt  and  a  large  amount  of  ammonium  chloride, 
the  precipitate  first  formed  was  redissolved  in  hydrochloric 
acid  and  precipitation  again  brought  about  by  the  very  careful 
addition  of  dilute  ammonia  in  slight  but  distinct  excess.  The 
mixture  was  heated  only  until  the  precipitate  became  silky 
and  crystalline,  when  it  was  allowed  to  stand  and  cool  for  a 
half  hour.  The  precipitate  was  filtered  off  upon  asbestos  in 
a  perforated  platinum  crucible  under  pressure,  ignited  and 
weighed.  Table  V  comprises  the  results  of  experiments  made 

*  Loc.  cit. 


AS   THE  PYROPHOSPHATE. 


131 


in  this  manner.  In  those  of  section  A  the  precipitation  was 
made  in  platinum  vessels ;  in  those  of  section  B  the  treatment 
was  in  glass. 


TABLE  V. 


A.    IN  PLATINUM. 

Mn2P207  equivalent 
to  MnO2. 

Error  in 
terms  of 
Mn,P207. 

Error  In 
terms  of 

Mil  11  £T<1  11686. 

Saturated 
solution  of 
HNaNH4P04. 

NH4C1. 

Total 
volume. 

Manganese 
in  the 
filtrate. 

Taken. 

Found. 

grm. 
0.1885 
0.1885 
0.1885 
0.1885 
0.3770 
0.3770 
0.3770 
0.3770 

gmi. 

0.1903 
0.1910 
0.1913 
0.1911 
0.3776 
0.3773 
0.3778 
0.3783 

grin. 

0.0018+ 
0.0025+ 
0.0028+ 
0.0026+ 
0.0006+ 
0.0003+ 
0.0008+ 
0.0013+ 

grm. 

0.0007+ 
0.0010+ 
0.0011+ 
0.0010+ 
0.0002+ 
0.0001+ 
0.0003+ 
0.0005+ 

cm8 
5 
5 
5 
5 
5 
5 
5 
5 

grm. 
20 
20 
20 
20 
20 
20 
20 
20 

cm» 
200 
200 
200 
200 
200 
200 
200 
200 

None. 
None. 
None. 
None. 
None. 
None. 
None. 
None. 

B.    IN  GLASS. 

0.1885 
0.1885 
0.3770 
0.3770 

0.1904 
0.1898 
0.3767 
0.3784 

0.0019+ 
0.0013+ 
0.0003- 
0.0014+ 

0.0007+ 
0.0005+ 
0.0001- 
0.0005+ 

6 
6 
5 
5 

20 
20 
20 
20 

200 
200 
200 
200 

None. 
None. 
None. 
None. 

In  this  series  of  experiments  the  mean  indication  is,  for  the 
first  time,  in  excess  of  the  theory.  Previously  the  error  has 
been  one  of  deficiency,  and  that  in  proportion  to  the  amount 
of  manganese  handled,  no  doubt  because  the  amount  of  uncon- 
verted trimanganese  phosphate  is  proportioned  to  the  entire 
amount  of  the  phosphate.  The  positive  error  which  is  devel- 
oped in  this  last  series  of  determinations  is  probably  due  to 
the  appearance  of  the  natural  error  of  all  precipitation  pro- 
cesses—  viz.,  the  tendency  on  the  part  of  the  precipitate  to 
include  matter  in  solution.  In  the  previous  experiments  this 
effect  was  doubtless  obscured  by  the  incompleteness  of  the 
conversion  of  the  trimanganous  phosphate  to  the  ammonium 
manganese  phosphate.  Indeed  it  is  quite  possible  that  even 
in  the  last  determinations  the  conversion  is  not  absolute,  and 
that  this  is  so  suggested  by  the  fact  that  the  errors  of  excess 
are  larger  in  the  case  of  the  smaller  amounts  of  manganese  for 


132  DETERMINATION  OF  MANGANESE 

which  the  conversion  throughout  the  entire  work  has  appeared 
to  be  more  complete.  From  the  consideration  of  the  results 
tabulated  and  described  it  would  seem  to  be  obvious  that  not 
only  is  the  presence  of  ammonium  chloride  not  objectionable 
in  this  analytical  process,  which  depends  upon  obtaining  the 
ammonium  manganese  phosphate  from  the  trimanganese  phos- 
phate precipitated  from  a  pure  solution  of  manganese,  but  that 
its  presence  in  not  too  small  amount,  or  that  of  a  substitute,  is 
absolutely  essential  to  make  this  conversion  complete.  For  a 
given  amount  of  manganese  and  a  given  volume  of  solution 
it  seems  essential  that  the  amount  of  ammonium  chloride 
should  reach  a  certain  limit.  According  to  our  experience 
the  proportion  of  ammonium  chloride  to  the  pyrophosphate 
should  be  at  least  50  : 1 ;  or,  speaking  approximately,  more 
than  200  molecules  of  ammonium  chloride  must  be  present  in 
the  liquid  (100  cm3  or  200  cm3)  to  every  molecule  of  the 
ammonium  manganese  phosphate  to  be  formed.  However, 
the  quantity  of  the  ammonium  salt  may  be  increased  almost 
to  the  point  of  saturation  of  the  liquid  without  causing  more 
than  a  trifling  solubility  of  the  ammonium  manganese  phos- 
phate in  the  presence  of  an  excess  of  the  precipitant.  The 
statement  of  Fresenius  and  Munroe  that  ammonium  manganese 
phosphate  is  soluble  in  ammonium  chloride  does  not  hold  if 
there  is  an  abundance  of  the  soluble  precipitating  phosphate 
present.  Further,  our  experience  goes  to  show  that  the  pre- 
cipitate may  be  washed  with  perfect  safety  with  pure  water  as 
well  as  with  slightly  ammoniacal  water,  or  with  ammoniacal 
water  containing  ammonium  nitrate,  if  the  nitration  is  per- 
formed rapidly  and  the  precipitate  is  gathered  in  small  space, 
as  is  the  case  when  the  phosphate  is  collected  on  asbestos  in 
a  perforated  crucible.  The  finely  granular  precipitate  which 
may  be  obtained  by  slow  action  of  dilute  ammonia  added 
gradually  to  the  hot  solution  of  the  manganese  salt  apparently 
includes  a  portion  of  unconverted  phosphate  which  resists  the 
replacement  of  the  manganese  by  ammonium.  On  the  other 
hand,  the  precipitate  of  flocky  condition  thrown  down  hi  the 
cold  passes  easily  to  the  silky  and  crystalline  condition  when 
heated  with  the  proper  amount  of  ammonium  salt  and  pos- 


AS   THE  PYROPHOSPHATE.  133 

sesses  a  constitution  approaching  the  ideal  under  such  condi- 
tions. The  conversion  of  the  flocky  manganous  phosphate  is 
so  rapid  that  the  precipitation  may  be  carried  on  safely  in 
glass  vessels.  If  the  ammonium  chloride  in  the  solution  were 
to  be  included  in  the  precipitate  it  would  volatilize  entirely 
during  the  ignition,  leaving  no  trace  unless,  possibly,  a  por- 
tion of  its  chlorine  were  to  combine  with  the  manganese. 
Tests  for  chlorine  in  the  residue  of  pyrophosphate  resulted 
negatively — no  more  than  a  mere  trace  being  found  in  any 
case,  so  that  the  contaminating  effect  of  the  ammonium  chlor- 
ide proves  to  be  insignificant  and  the  responsibility  for  the 
increase  in  weight  above  the  theory  must  apparently  rest  with 
the  included  microcosmic  salt. 

In  the  practical  determination  of  manganese  by  the  phos- 
phate method  of  Gibbs,  therefore,  we  advocate  strongly  the 
presence  of  large  amounts  of  ammonium  chloride.  Good 
results  may  be  obtained  by  the  method  of  precipitation  origi- 
nally laid  down  by  Gibbs,  or  by  the  modification  proposed 
by  Blair,  if  the  ammonium  salt  is  present  in  sufficient  quantity. 
On  the  whole  trustworthy  results  are  obtained  most  easily  and 
surely,  according  to  our  experience,  by  following  the  method 
of  the  experiments  of  Table  V.  The  slightly  acid  solution, 
containing  in  a  volume  of  200  cm3  (in  platinum  or  glass)  an 
amount  of  manganese  not  more  than  enough  to  make  0.4  grm. 
of  the  pyrophosphate,  20  grm.  of  ammonium  chloride  and  5 
to  10  cm3  of  a  cold  saturated  solution  of  microcosmic  salt,  is 
precipitated  in  the  cold  by  the  careful  addition  of  dilute 
ammonia  in  only  slight  excess.  The  mixture  is  heated  until 
the  precipitate  becomes  silky  and  crystalline,  the  whole  is 
allowed  to  stand  and  cool  a  half  hour,  the  precipitate  is  col- 
lected upon  asbestos  in  a  perforated  platinum  crucible,  washed 
(best  with  slightly  ammoniacal  water),  dried  at  gentle  heat 
and  ignited  as  usual.  By  this  process  determinations  of  the 
larger  amounts  of  manganese  —  0.4  grm.  of  the  pyrophosphate 
—  approximate  rather  more  closely  to  the  theoretical  values 
than  do  those  of  the  smaller  amounts  —  0.15  grm.  In  either 
case  the  average  error  should  not  exceed  0.0010  grm.  in  terms 
of  manganese. 


XVIII 

ON  THE  DETECTION  OF  SULPHIDES,  SULPH- 
ATES, SULPHITES  AND  THIOSULPHATES  IN 
THE  PRESENCE  OF  EACH  OTHER. 

BY  PHILIP  E.  BROWNING  AND  ERNEST  HOWE  * 

SOME  three  years  ago  R.  Greig  Smith  f  published  a  method 
for  the  detection  of  sulphates,  sulphites  and  thiosulphates  in 
the  presence  of  each  other,  which  promised  much  toward  the 
solution  of  this  most  difficult  problem.  The  method  may  best 
be  described  in  the  author's  own  language :  To  a  solution  of 
the  salts  of  the  above  mentioned  acids  "barium  chloride  is 
added  in  excess,  together  with  a  good  quantity  of  ammonium 
chloride,  which,  like  many  salts  of  ammonium,  potassium  and 
calcium,  acts  as  a  flocculent  or  coagulant,  and  facilitates  the 
filtration  of  the  barium  sulphate.  Hydrochloric  acid  is  next 
added,  drop  by  drop,  until  it  is  evident  that  there  is  no  further 
solution  of  barium  sulphite  and  thiosulphate,  and  that  only  the 
sulphate  remains  undissolved;  the  solution  is  then  filtered 
through  a  moistened  double  filter  paper,  which  should  be  free 
from  'pin  holes.'  The  filtrate  will  probably  be  clear,  but  if 
not  it  should  be  returned  to  the  filter  for  a  second  filtration. 
When  too  much  thiosulphuric  acid  is  present,  the  clear  filtrate 
will  visibly  become  clouded,  or  from  being  whitish  will  become 
more  opaque;  if  this  occurs  the  solution  should  be  thrown 
out,  and  a  fresh  portion  made  more  dilute.  A  solution  of 
iodine  is  added  to  half  of  the  filtrate  until  the  color  is  of  a 
permanent  yellow  tinge ;  a  white  precipitate  indicates  the 
presence  of  a  sulphite  which  has  been  oxidized  by  the  iodine 

*  From  Am.  Jour.  Sci.,  vi,  317.  t  Chem.  News,  Ixxii,  39. 


DETECTION  OF  SULPHIDES,  SULPHATES,  ETC.      135 

to  sulphate.  In  the  absence  of  a  decided  precipitate  traces  of 
sulphite  may  easily  be  detected  by  comparing  the  treated  and 
untreated  halves  of  the  filtrate  —  a  procedure  which  very  often 
saves  a  good  deal  of  time,  as  it  is  unnecessary  to  wait  until  a 
clear  filtrate  is  obtained.  The  two  halves  are  mixed,  and  if 
the  yellow  color  disappears  more  iodine  is  added,  the  solution 
filtered  and  the  filtrate  divided  into  two  halves  as  before. 
With  a  slight  turbidity  filtration  may  be  omitted.  Bromine 
water  is  added  to  one  of  the  halves  when  any  thiosulphate  in 
the  original  solution  shows  itself  as  a  white  precipitate  of 
barium  sulphate,  readily  seen  on  comparing  the  two  test-tubes. 
The  thiosulphate  is  by  iodine  converted  to  tetrathionate,  which 
is  oxidized  by  bromine  water  to  sulphate."  Three  objections 
to  this  method  as  described  will  readily  occur  to  the  reader : 
first,  the  readiness  with  which  the  thiosulphate  is  decomposed 
by  free  hydrochloric  acid;  second,  the  comparatively  large 
amount  of  acid  necessary  to  effect  the  complete  solution  of  the 
barium  sulphite  and  thiosulphate  when  precipitated  with  the 
sulphate  as  compared  with  the  amount  required  to  prevent 
the  precipitation ;  third,  the  lack  of  delicacy  necessitated  by  a 
comparison  of  portions  of  a  colored  solution  in  looking  for 
small  precipitates.  The  work  to  be  described  was  undertaken 
to  overcome  these  difficulties  and  to  test  the  accuracy  of  a 
modified  method.  Solutions  of  potassium  sulphite  and  sodium 
thiosulphate  were  made  approximately  decinorrnal  and  stand- 
ardized in  the  usual  manner  against  an  iodine  solution  of 
known  value.  It  was  found  that  by  making  a  solution 
containing  sulphates,  sulphites,  and  thiosulphates  very  faintly 
acid,  the  sulphites  and  thiosulphates  were  held  completely  in 
solution  when  the  barium  sulphate  was  precipitated.  The 
extreme  sensitiveness  of  a  thiosulphate  to  the  decomposing 
action  of  free  hydrochloric  acid  suggested  the  possible  substi- 
tution of  acetic  acid  to  hold  the  sulphites  and  thiosulphates 
in  solution.  This  being  a  weaker  acid,  we  hoped  to  avoid  the 
decomposition  of  the  thiosulphate  into  sulphur  and  sulphurous 
acid,  or  at  least  to  delay  the  decomposing  action.  The  results 
of  these  experiments  appear  in  the  following  table : 


136      DETECTION  OF  SULPHIDES,  SULPHATES,  ETC., 

TABLE  I. 


TTWI,,  «,  a 

Hydro- 

E 

bcp. 

\  oiume 
cm3  of 
water. 

chloric 
acid 
(1:4). 

Acetic 
acid. 

NaAO, 
taken. 

Result. 

drops. 

drops. 

grin. 

• 

1) 

10 

2 

.  . 

0.01 

No  sulphur  in  20  minutes. 

10 

2 

.  . 

0.1 

Sulphur  in  45  seconds. 

3) 

100 

3 

.  . 

0.1 

Sulphur  in  15  minutes. 

4) 

10 

. 

8 

0.01 

No  sulphur  in  20  minutes. 

5) 

10 

. 

8 

0.1 

Sulphur  in  90  seconds. 

(6) 

100 

, 

10 

0.1 

No  sulphur  in  20  minutes. 

(7) 

100 

, 

10 

0.25 

Sulphur  in  15  minutes. 

i 

(8) 
(9 

100 
100 

•' 

10 
10 

0.6 
1.0 

Sulphur  in  60  seconds. 
Sulphur  in  30  seconds. 

From  these  results  it  would  seem  that  the  decomposition  of 
a  thiosulphate  is  more  rapid  in  presence  of  hydrochloric  acid 
than  hi  presence  of  a  much  larger  amount  of  acetic  acid. 

Our  next  experiments  were  directed  toward  a  determination 
of  the  effect  of  adding  stannous  chloride  to  bleach  the  color 
of  the  free  iodine  and  bromine  used  in  the  oxidation  and  of 
acidifying  with  acetic  acid,  before  treating  with  barium 
chloride.  That  is  to  say,  the  process  as  we  used  it,  consisted 
in  acidifying  the  solution  to  be  tested  with  acetic  acid,  adding 
barium  chloride,  filtering  to  remove  precipitated  sulphate 
(always  present  in  the  sulphite),  adding  iodine  to  the  nitrate 
until  the  color  was  permanent,  bleaching  with  stannous 
chloride,  filtering  off  the  sulphate  which  represents  the  sul- 
phite originally  present,  adding  bromine  in  excess  to  the 
filtrate  and  again  bleaching  with  stannous  chloride  to  increase 
the  visibility  of  the  sulphate  which  now  represents  the 
thiosulphate  originally  present.  The  details  of  experiments 
in  which  the  sulphite  was  taken  alone  and  oxidized  with 
iodine  are  given  in  Table  II. 

A  corresponding  series  of  experiments  was  made  in  which 
hydrochloric  acid  was  substituted  for  acetic  acid  and  essentially 
the  same  results  were  obtained. 

A  similar  series  of  experiments  was  made  to  test  the  effect 
of  treating  the  thiosulphate  in  an  acidified  solution,  first  with 


IN  THE  PRESENCE  OF  EACH  OTHER. 


137 


TABLE  H. 


Brp. 

KJO. 

taken. 

Volume 
of 
water. 

BaS04  precipitated 
after  oxidation 
with  iodine. 

Remarks. 

(1) 
(2) 
(3) 
(4) 
(5) 

gnu* 

0.1 
0.01 
0.001 
0.0005 
0.0001 

cm8 
10 
10 

10 
10 
10 

Very  abundant. 
Abundant. 
Distinct. 
Fair. 
Faint. 

Plainly  visible  before  adding  SnCl2. 
Plainly  visible  before  adding  SnCl2. 
More  distinct  after  adding  SnCl2. 
Hardly  visible  before  adding  SnCl2. 
Invisible  before  adding  SnCLj. 

iodine  and  then  after  filtration  (if  a  precipitate  had  formed) 
with  bromine.  In  the  experiments  of  division  A  hydrochloric 
acid  (a  few  drops)  was  added  before  treating  with  barium 
chloride,  and  in  those  of  division  B  acetic  acid  was  used 
similarly.  Stannous  chloride  was  employed  to  bleach  the 
excess  of  iodine  and  bromine. 

TABLE  m. 


BaS04  pre- 

BaSO4 precip- 

Exp. 

Na2S2Os 
taken. 

Volume 
of  water. 

cipitated 
by  action 

itated  by 
action 

Remarks. 

of  iodine. 

of  bromine. 

A. 

grm. 

cm3 

(1) 

0.1 

10 

Distinct. 

Abundant. 

Sulphur  separated  in  30  seconds. 

(2) 

0.01 

10 

Faint. 

Abundant. 

No  sulphur  in  90  seconds. 

(3) 

0.001 

10 

None. 

Distinct. 

No  sulphur  in  several  minutes. 

(4) 

0.0005 

10 

None. 

Faint. 

No  sulphur  ;  SnCl2  necessary. 

(5) 

0.0001 

10 

None. 

Very  faint. 

No  sulphur  ;  SnCl2  necessary. 

B. 

(1) 

0.1 

10 

Faint. 

Abundant. 

No  sulphur  separated  in  1  minute. 

(2) 

0.01 

10 

None. 

Abundant. 

No  sulphur  separated  in  several 

minutes. 

(3) 

0.001 

10 

None. 

Distinct. 

No  sulphur. 

(4) 

0.0005 

10 

None. 

Faint. 

No  sulphur  ;  SnCl2  necessary. 

(5) 

0.0001 

10 

None. 

Very  faint. 

No  sulphur  ;  SnCl2  necessary. 

From  these  experiments  the  advantage  of  the  use  of  acetic 
acid  becomes  apparent,  as  does  also  the  use  of  stannous 
chloride  in  increasing  the  delicacy  of  this  indication,  so  that  a 
small  fraction  of  a  milligram  may  easily  be  detected. 


138      DETECTION  OF  SULPHIDES,  SULPHATES,  ETC., 


If  relatively  large  amounts  of  thiosulphate  are  present 
with  small  amounts  of  sulphite,  we  have  sometimes  found  it 
advantageous  to  manipulate  so  that  even  the  slow  decom- 
position of  the  thiosulphate  by  acetic  acid  may  be  avoided 
by  first  attempting  precipitation  with  barium  chloride  in  a 
dilute  ammoniacal  solution.  By  this  method  the  barium 
sulphate  and  sulphite  are  separated  from  the  thiosulphate 
and  identified — the  sulphate  by  its  insolubility  in  dilute 
hydrochloric  acid,  and  the  sulphite  by  the  action  of  iodine 
upon  the  acid  filtrate  from  the  barium  sulphate.  After 
filtering,  the  thiosulphate  may  be  detected  in  the  filtrate  by 
the  use  of  iodine  and  bromine  as  described  above.  Table  IV 
gives  some  results  by  this  treatment. 

TABLE  IV. 


Exp. 

Na2S308 
taken. 

BaS04  pre- 
cipitated by 
iodine. 

BaS04  pre- 
cipitated by 
bromine. 

Remarks. 

grm. 

(1) 

0.1 

None. 

Abundant. 

(2) 

0.01 

None. 

Good. 

(3) 

0.001 

None. 

Fair. 

SnCl2  necessary. 

(4) 
(5) 

0.0005 
0.0001 

None. 
None. 

Faint. 
None. 

SnCl2  necessary. 

As  will  be  seen,  the  test  for  the  thiosulphate  by  this  method 
of  treatment  is  not  so  delicate,  probably  on  account  of 
mechanical  holding  of  the  barium  thiosulphate  by  the  pre- 
cipitated sulphate  and  sulphite. 

Having  determined  the  limits  of  accuracy  of  the  method  as 
applied  to  the  sulphite  and  thiosulphate  taken  separately,  our 
next  experiments  were  directed  toward  an  investigation  of  the 
working  of  the  method  when  these  two  acids  are  found 
together  in  solution.  Sulphates,  almost  invariably  present 
with  sulphites,  are  of  course  quite  easily  separated  by  filtration 
and  treating  with  the  barium  salt  in  acid  solution.  Sulphides 
if  present  in  the  solution  would  seriously  interfere  with  the 
working  of  this  method  if  not  removed,  being  readily  oxidized 
by  the  iodine  or  bromine  to  sulphite,  sulphate,  or,  should 


IN  THE  PRESENCE  OF  EACH  OTHER.     139 

sulphur  also  separate,  to  thiosulphate.  We  found  in  course 
of  our  work  that  in  attempting  to  neutralize  a  mixture  of 
freshly  prepared  alkaline  sulphide  together  with  a  sulphite 
we  often  obtained  a  precipitate  of  sulphur.  After  the  removal 
of  the  sulphide  and  sulphate,  we  were  surprised  to  find  on 
treating  with  iodine  scarcely  a  trace  of  sulphite.  On  treating 
with  bromine  however  an  abundant  indication  of  thiosulphate 
was  obtained.  It  is  well  known  of  course  that  thiosulphate 
may  be  formed  by  boiling  a  sulphite  with  sulphur,  but  that 
this  reaction  should  take  place  so  readily  and  completely 
seemed  to  us  rather  unusual. 

For  the  removal  of  a  sulphide  before  proceeding  with  the 
tests  for  sulphite  and  thiosulphate  Greig  Smith  recommends 
the  passing  of  carbon  dioxide  through  the  solution  until  the 
escaping  gas  gives  no  indication  of  hydrogen  sulphide,  but 
Bloxam*  calls  attention  to  the  tedious  and  wholly  unsatis- 
factory character  of  this  method  of  removal  and  recommends 
a  mixture  of  zinc  chloride,  cadmium  chloride,  ammonium 
chloride  and  ammonia.  We  have  found  that  the  addition  of 
zinc  acetate  to  a  faintly  alkaline  solution  accomplishes  the 
same  purpose  in  an  entirely  satisfactory  manner.  The  sul- 
phide used  in  our  work  was  freshly  made  by  passing  hydrogen 
sulphide  through  a  dilute  solution  of  sodium  hydroxide. 
When  portions  of  this  solution,  still  alkaline,  were  treated 
with  zinc  acetate  in  excess,  and  the  zinc  hydroxide  and 
sulphide  removed  by  filtration,  the  filtrate  gave  no  test  for 
either  sulphite  or  thiosulphate  by  the  application  of  iodine 
and  bromine  as  described,  and  the  vapor  evolved  on  boiling 
caused  no  darkening  of  lead  paper.  The  following  table 
shows  the  results  of  a  few  experiments  in  which  tests  were 
made  for  the  sulphite  and  thiosulphate,  after  removing  a  con- 
siderable amount  of  the  sulphide  in  the  manner  described,  and 
of  the  sulphate  by  acidifying  and  adding  barium  chloride. 

The  method  as  we  have  modified  it  may  be  summarized  as 
follows :  To  about  0.1  grm.  of  the  substance  to  be  analyzed 
dissolved  in  10  cm8  of  water  or  more,  add  sodium,  potassium  or 
*  Chem.  News,  Ixxii,  63. 


140      DETECTION  OF  SULPHIDES,   SULPHATES,  ETC. 

TABLE  V. 


E2SO3  taken. 

N^SjOj,  taken. 

BaS04  precipitated 
after  oxidation 
with  iodine. 

BaS04  precipitated 
after  oxidation 
with  bromine. 

grin* 

0.1 
0.1 

0.01 
0.001 
0.001 

grin. 

0.01 
0.001 
0.1 
0.1 
0.001 

Abundant. 
Abundant. 
Good. 
Faint. 
Fair. 

Good. 
Distinct. 
Abundant. 
Abundant. 
Fair. 

ammonium  hydroxide  to  distinct  but  faintly  alkaline  reaction. 
The  solution  should  be  neutral  or  alkaline  rather  than  even 
faintly  acid,  owing  to  the  readiness  with  which  sulphur 
separates.  To  the  alkaline  solution  add  zinc  acetate  in 
distinct  excess  and  filter.  The  precipitate  may  be  tested  for 
hydrogen  sulphide,  on  acidifying,  in  the  usual  manner.  To 
the  filtrate  add  acetic  acid,  a  few  drops  in  excess  of  the 
amount  necessary  to  neutralize,  and  barium  chloride,  and  filter 
through  a  double  filter.  To  the  filtrate  add  iodine  until  the 
solution  takes  on  a  permanent  yellow  tinge,  and  then  bleach 
with  stannous  chloride,  best  after  adding  a  few  drops  of 
hydrochloric  acid  to  prevent  the  possible  precipitation  of  a 
basic  salt  of  tin.  A  precipitate  at  this  point  indicates  the 
sulphite.  Filter,  add  bromine  water  in  faint  excess  to  the 
filtrate,  bleaching  again  with  stannous  chloride.  A  pre- 
cipitate on  adding  bromine  indicates  a  thiosulphate  originally 
present. 


XIX 

ON  THE   SEPARATION   OF   NICKEL  AND 
COBALT  BY  HYDROCHLORIC  ACID. 

BYFRANKE  STUART  HAVENS. 

A  QUANTITATIVE  separation  of  nickel  and  cobalt  by  a  process 
analogous  to  that  published  from  this  laboratory  for  the 
separation  of  aluminum  and  ironf  has  been  put  forward  in  a 
recent  paper  by  E.  Pinerua.J  The  process  may  be  described 
briefly  as  follows:  The  hydrous  chlorides  of  nickel  and 
cobalt  (0.3-0.4  grms.)  are  dissolved  in  a  little  water  and  to  the 
solution  are  added  10  to  12  cm3  of  aqueous  hydrochloric  acid 
and  10  cm3  of  ether,  and  the  whole,  contained  in  a  little 
beaker  surrounded  with  water,  and  ice,  is  saturated  with 
gaseous  hydrochloric  acid.  The  cobalt,  which  remains  in 
solution,  is  decanted  off  and  the  yellow  insoluble  nickel 
chloride  washed  with  a  previously  prepared  solution  of  ether 
saturated  with  hydrochloric  acid  gas  at  a  low  temperature. 
The  nickel  is  determined  by  known  methods,  preferably  as 
the  sulphate.  The  author  claims  very  precise  results  for  the 
process,  but  gives  no  experimental  proof  of  his  work.  Previous 
to  the  appearance  of  this  paper  my  experiments  upon  the 
solubility  of  nickel  chloride  in  an  ether-hydrochloric  acid 
solution,  such  as  used  in  our  process  for  the  separation  of 
aluminum  and  iron,  which  is  practically  the  same  in  proportions 
as  that  used  by  Pinerua  to  effect  precipitation,  had  shown  that, 
while  nickel  chloride  is  somewhat  insoluble  in  such  a  mixture, 
the  degree  of  insolubility  is  not  sufficient  for  a  quantitative 
separation.  Since  the  appearance  of  Pinerua's  work  I  have 

*  From  Am.  Jour.  Sci.,  vi,  396. 

t  Gooch  and  Havens,  Am.  Jour.  Sci.,  ii,  416.    This  volume,  p.  20. 

t  Gaz.  chim.  ital.,  xxvii,  56. 


142  ON  THE  SEPARATION  OF  NICKEL  AND 

been  over  the  ground  again  and  have  reached  the  same 
conclusions  as  before,  as  shown  in  the  following  experiments. 

When  a  solution  of  0.02  grm.  of  nickel  chloride  (free  from 
iron  and  cobalt)  in  7  cm3  of  aqueous  hydrochloric  acid,  was 
saturated  with  hydrochloric  acid  gas  at  a  temperature  of  —2° 
C.  (obtained  by  immersing  the  container  in  a  mixture  of  ice 
and  salt)  no  precipitation  resulted.  When,  however,  an  equal 
volume  of  ether  was  added  and  the  whole  was  again  saturated 
with  hydrochloric  acid  gas  a  yellow  precipitate  formed,  while 
the  supernatant  liquid  still  remained  of  a  deep  green  color. 
The  solution  was  filtered  quickly  through  asbestos  in  a  filter 
crucible,  and  the  clear  filtrate  after  evaporation  with  sulphuric 
acid  was  electrolyzed.  The  metallic  deposit  of  0.0020  grm. 
proved  to  be  pure  nickel ;  for  when  dissolved  in  nitric  acid  it 
gave  no  test  for  iron  with  potassium  sulphocyanide  or  ferro- 
cyanide,  and  neither  the  apple-green  hydroxide  nor  the  black 
sulphide,  prepared  by  the  usual  methods,  showed  any  trace  of 
cobalt  in  the  borax  bead.  It  is  obvious,  therefore,  that  nickel 
chloride  is  not  fully  precipitated  under  these  conditions  and 
that  the  green  color  of  the  solution  is  due  to  nickel  in  solution 
and  not  to  traces  of  iron,  as  Pinerua  has  supposed.*  A  second 
experiment  similar  to  the  first  showed  a  solubility  of  the 
nickel  chloride  represented  by  0.0018  grm.  of  metallic  nickel. 
It  is  evident,  then,  that  the  solubility  of  nickel  chloride  in  this 
mixture  of  aqueous  hydrochloric  acid  and  ether  thoroughly 
saturated  with  hydrochloric  acid  gas  is  not  far  from  an  amount 
represented  by  0.0020  grm.  of  metallic  nickel  for  every  14  cm8 
of  solution. 

Still  another  experiment,  in  which  nickel  chloride  repre- 
senting 0.0020  grm.  of  metallic  nickel  was  treated  with  14  cm3 
of  the  ether-hydrochloric  acid  solution  and  the  whole  saturated 
for  one  hour  at  a  low  temperature  with  hydrochloric  acid  gas 
without  precipitation,  showed  the  same  thing. 

When  the  nickel  chloride  remaining  on  the  asbestos  was 
washed  with  about  40  cm8  of  a  mixture  of  equal  parts  ether  and 
aqueous  hydrochloric  acid  saturated  with  hydrochloric  acid  gas, 

*  Loc.  cit.  ' 


COBALT  BY  HYDROCHLORIC  ACID.  143 

the  washings  evaporated  with  sulphuric  acid  and  treated  by 
the  battery  gave  a  deposit  of  metallic  nickel  weighing  0.0027 
grm.  —  an  amount  proportionately  less  than  that  found  in  the 
nitrate  proper. 

Although  employing  a  mixture  of  aqueous  hydrochloric  acid 
and  ether  saturated  with  gaseous  hydrochloric  acid  for  the 
precipitation,  Pineriia  has  advised  the  use  of  pure  ether 
saturated  with  gaseous  hydrochloric  acid  for  the  washing.  In 
my  experiments  with  such  a  mixture  I  find  that  in  it  the 
hydrous  nickel  chloride  is  practically  insoluble  and  that  30  cm3 
of  the  washings  of  the  precipitated  chloride  with  such  a 
mixture  gave  no  deposit  of  nickel  by  the  battery.  It  seemed 
possible,  therefore,  that  by  reducing  the  water  present  to  the 
lowest  possible  amount  necessary  to  dissolve  the  chlorides  to 
be  treated  the  precipitation  of  the  nickel  might  be  made  more 
complete.  The  experiments  of  the  following  table  were  made 
to  put  this  idea  to  the  test. 

Solutions  of  the  pure  chlorides  of  nickel  and  cobalt,  carefully 
purified  and  freed  from  other  metals  and  each  other,  were, 
after  conversion  to  the  form  of  the  sulphate,  standardized 
by  the  battery.  Weighed  portions  of  these  solutions  were 
taken  in  a  small  beaker,  evaporated  to  dryness,  the  dry  salts 
dissolved  in  as  little  water  as  possible  (about  1  cm3),  10  to  15  cm3 
of  ether  added,  and  the  whole  saturated  with  hydrochloric  acid 
gas,  the  beaker  being  meanwhile  immersed  in  running  water 
and  cooled  to  about  15°  C.  When  saturation  was  complete 
the  precipitated  chloride  was  caught  on  asbestos  in  a  filter 
crucible,  washed  thoroughly  with  a  previously  saturated 
solution  of  ether,  dissolved  in  water,  evaporated  with  sulphuric 
acid  and  determined  as  metallic  nickel  by  the  battery.  The 
cobalt  in  the  filtrate  was  recovered  by  evaporation  and 
electrolysis  in  like  manner. 

Experiments  (1),  (2),  and  (3)  of  the  accompanying  table 
show  that  by  this  process  the  nickel  is  thrown  down  quantita- 
tively, and  experiments  (2)  and  (3)  show  that  in  the  presence 
of  a  few  milligrams  of  the  cobalt  salt  the  separation  of  a  small 
amount  of  nickel  is  sharp.  The  residue  of  nickel  in  these 


144 


SEPARATION  OF  NICKEL  AND   COBALT. 


experiments  gave  no  test  for  cobalt  with  the  borax  bead. 
When,  however,  the  cobalt  is  present  to  the  amount  of  a  few 
centigrams  as  in  (4),  (5),  and  (6),  the  precipitated  nickel 
chloride,  which  forms  a  hard  mass,  includes  the  cobalt  salt  so 
that  even  a  large  quantity  of  washing  solution  (100  cm3  was 
used  in  experiment  6)  cannot  remove  it. 


Exp. 

Nickel  taken 
as  the 
hydrous 
chloride. 

Nickel 
found. 

Error. 

Cobalt  taken 
as  the 
hydrous 
chloride. 

Cobalt 
found. 

Error. 

(1) 

(2) 
(3) 
(4) 
(5) 
(6) 

giro. 

0.0068 
0.0090 
0.0090 
0.0469 
0.0468 
0.0472 

grlu. 

0.0066 
0.0090 
0.0091 
0.0490 
0.0503 
0.0493 

grm. 
0.0002- 

0.0000 
0.0001+ 
0.0021+ 
0.0035+ 
0.0021+ 

grm. 

0.0030 
0.0123 
0.0700 
0.0700 
0.0700 

grin. 

0.0127 

grin. 

0.0004+ 

From  the  experiments  described  it  is  obvious  that  the  pro- 
cess as  proposed  by  Pinerua  will  not  give  a  complete  precipi- 
tation of  the  nickel  chloride.  Nickel  chloride  is,  however, 
practically  insoluble  in  pure  ether  saturated  with  hydrochloric 
acid  gas  and  can  be  separated  from  small  quantities  of  the  solu- 
ble cobalt  salt  in  that  medium.  In  the  presence  of  even  a  few 
centigrams  of  the  cobalt  chloride,  however,  the  process  is  not 
practicable  on  account  of  the  inclusion  of  the  cobalt  by  the 
massive  nickel  chloride.  It  is  possible  that  by  repeated  solu- 
tions and  reprecipitations  the  nickel  salt  might  be  sufficiently 
freed  from  the  cobalt,  but  the  process  must  naturally  be  long 
and  tedious. 


XX 

THE    ETHERS    OF    TOLUQUINONEOXIME    AND 

THEIR  BEARING  ON  THE  SPACE  ISOMER- 

ISM  OF  NITROGEN. 

BY  JOHN  L.  BRIDGE  AND  WILLIAM  CONGER  MORGAN  * 

IN  an  article  on  the  ethers  of  quinoneoxine  (isonitrosophenol) 
published  by  one  of  us,f  it  was  stated,  that  when  boiled  with 
alcohol,  the  benzoyl  ether  of  quinoneoxime  dichloride  gave 
gave  two  monochlor  substitution-products.  Naturally  it  was 
supposed  that  the  chlorine  atom  occupied,  in  the  one,  an  ortho 
position,  and,  in  the  other,  a  meta  position  to  the  radical 
NOR,  the  reaction  being : 

000 


+  +  2HC1. 

Hk       /"HC1      HL         JH      Hi         JCl 


N  N  N 

O  00 

R  R  R 

It  was  also  found  that  these  same  isomers  were  formed  when 
monochlorquinone  was  treated  with  hydroxylamine  hydro- 
chloride,  and  the  sodium  salt  of  the  chlorquinoneoxime  thus 
formed,  treated  with  benzoyl  chloride. 

The  preceding  explanation  regarding  the  splitting  off  of 
hydrochloric  acid  from  the  dichloride  is  not  in  accord  with  re- 
sults of  work  done  by  Kehrmann,J  published  in  his  article  on 
the  influence  of  radicals  in  the  action  of  hydroxylamine  on 

*  From  Am.  Chem.  Jour.,  xx,  761. 
t  Ann.  Chem.  (Liebig),  cclxvii,  79. 
t  Ber.  Dtsch.  chem.  Ges.,  xxi,  3315 ;  Jour,  prakt.  Chem.  [2],  xl,  268. 

VOL.    II.  —  10 


146  THE  ETHERS   OF  TOLUQUINONEOXIME. 

quinones,  in  which  he  generalizes  the  results  of  his  observa- 
tions in  the  statement  that  the  presence  of  a  radical  attached 
to  the  ring  so  much  lessens  the  replaceability  of  the  quinone 
oxygen  atom  neighboring  to  the  radical  that  the  principal 
part  if  not  the  whole  of  the  resulting  product,  is  a  meta-sub- 
stituted  quinoneoxime.  The  work  of  his  former  article  has 
been  repeated  by  Kehrmann,*  who  finds  the  same  isomeric 
benzoyl  ethers  of  monochlorquinoneoxime ;  but  believing 
that  observations  of  their  behavior  indicate  the  substantiation 
of  his  rule,  he  states  that  both  these  ethers  have  the  chlorine 
atom  in  the  meta  position  to  the  oximido  group,  and  attributes 
their  difference  to  space  isomerism  of  nitrogen,  writing  the  re- 
action thus : 


+  2HC1. 


Kehrmann  designates  the  compound  represented  by  symbol  I.  as 
"  chlorquinonemetaantioxime  ether  "  and  II.  as  "  chlorquinone- 
metasynoxime  ether."  The  question  as  to  whether  the  chlor- 
ine atom  occupies  the  same  position  in  both  compounds  could 
be  definitely  settled  if  ortho-  and  metachlorphenols  could  be 
changed  to  the  corresponding  chlorquinoneoximes  or  so-called 
isonitrosophenols.  This  transformation  has,  however,  unfor- 
tunately not  yet  been  accomplished,  and  repeated  efforts  to 
obtain  the  corresponding  bromquinoneoximes  have  resulted 
in  failure,  orthobromphenol  not  being  attacked  by  nitrous 
acid  or  amyl  nitrite. 

The  toluquinoneoximes,  obtained  by  the  action  of  nitrous 
acid  on  ortho-  and  metacresol,  we  have  taken  up  for  study  as 
being  the  most  closely  analogous  compounds  in  which  the 
position  of  the  side  groups  is  definitely  known,  believing  that 

*  Ann.  Chem.  (Liebig),  cclxxix,  27. 


THE  ETHERS  OF  TOLUQUINONEOXIME.  147 

by  an  examination  of  these  bodies,  light  may  be  thrown  upon 
the  nature  of  the  others. 

When  orthocresol  is  acted  upon  by  nitrous  acid,  a  toluqui- 
nonemetaoxine  is  formed  according  to  the  following  reaction: 

H 

O 


+  HONO  = 
H 


H 


The  metacresol  forms  similarly  a  corresponding  orthooxime : 

H 

O  O 

HrXXxiH 

+  HONO  ==  +  H20. 

H  LJ  CH8 

H  N 

O 
H 

The  benzoyl  ethers  of  the  toluquinoneoximes  prove  to  be 
exceedingly  well  fitted  to  characterize  these  bodies  since 
they  are  formed  from  the  oximes  in  quantitative  proportions, 
are  easily  crystallized,  and  readily  distinguished  from  each 
other. 

The  benzoyl  ether  of  toluquinoneorthooxime  produced 
from  metacresol,  crystallizes  in  light  brownish-yellow  crystals 
melting  at  177°  C.  and  appearing  under  the  microscope  as 
long  rectangular  blades,  suggesting  the  orthorhombic  system. 
The  benzoyl  ether  of  toluquinonemetaoxime  produced  from 
orthocresol  is  obtained  in  the  form  of  yellow  crystals  which 
begin  to  soften  at  about  150°  and  do  not  melt  entirely 
until  at  about  190°.  Subjected  to  fractional  crystallization, 
a  portion  of  the  metaoxime  is  readily  obtained  consisting  of 
branching  needle-like  crystals,  melting  at  193°  C.  These,  as 


148  THE  ETHERS  OF  TOLUQUINONEOXIME. 

well  as  lower-melting  fractions,  can  be  readily  distinguished 
from  the  orthooxime  by  their  less  regular  appearance  under 
the  lens,  as  well  as  by  their  crystal  habit,  tending  to  produce 
curved  forms. 

The  marked  difference  in  the  form  and  habit  of  the  crystals 
of  the  benzoyl  ethers  of  the  ortho-  and  metatoluquinone 
oximes  makes  it  possible  to  study  carefully  the  product  of  the 
action  of  toluquinone  on  hydroxylamine.  Goldschmidt  and 
Schmidt  *  have  shown  that  the  principal  product  of  this  reac- 
tion is  toluquinonemetaoxime.  This  they  demonstrated  by 
oxidizing  to  a  dinitro  compound;  but  as  some  of  the  theo- 
retically possible  nitro-derivatives  of  the  cresols  are  not 
known,  this  method  can  scarcely  be  considered  to  prove  con- 
clusively the  absence  of  toluquinoneorthooxime.  We  have, 
therefore,  studied  further  the  product  of  the  action  of  hydroxyl- 
amine on  toluquinone  by  the  aid  of  the  benzoyl  ethers.  When 
the  sodium  salt  of  the  oxime  thus  formed  was  treated  in  alco- 
holic solution  with  benzoyl  chloride,  and  the  benzoyl  ether 
obtained  was  subjected  to  fractional  crystallization,  not  a 
trace  of  the  orthooxime  ether,  crystallizing  in  blades  and 
melting  at  177°,  was  found.  This  proves  that  when  tolu- 
quinone is  treated  with  hydroxylamine  the  whole  of  the  re- 
sulting product  is  toluquinonemetaoxime.  The  inference  is 
plain  that  Kehrmann's  rule  concerning  the  influence  of  side- 
chains,  attached  to  the  ring  in  quinones,  upon  the  entering 
oximido  radical  is  quantitatively  true  in  this  case ;  and  similar 
indications  furnished  by  the  corresponding  methyl  ethers,  as 
will  appear  later,  strengthen  this  conclusion. 

As  has  been  previously  shown,  the  benzoyl  ether  of  the 
metaoxime  does  not  consist  of  a  single  compound,  but  is  a 
mixture.  The  body  melting  at  193°  is  readily  separated  in 
considerable  quantities,  while  the  remainder  consists  of  a 
very  intimate  mixture  of  this  body  with  one  of  much  lower 
melting-point,  separated,  if  it  can  be  separated  at  all,  only 
with  the  greatest  difficulty.  By  concentration  of  the  mother- 
liquor  from  which  the  high-melting  fractions  have  been 

*  Ber.  Dtsch.  chem.  Ges.,  xvii,  2063. 


THE  ETHERS  OF  TOLUQUINONEOXIME.  149 

obtained,  and  carefully  crystallizing,  fractions  may  be  separated 
melting  almost  completely  at  temperatures,  varying  with  the 
fraction,  between  130°  and  150°.  Repeated  crystallizations 
separated  each  fraction  into  portions  melting,  on  the  one  hand, 
always  considerably  higher,  and,  on  the  other  hand,  often 
somewhat  lower,  and  ordinarily  no  definite  body  of  low 
melting-point  could  be  isolated.  Twice,  however,  a  nearly 
complete  separation  seemed  to  have  been  obtained.  Thus, 
from  fractions  melting  between  140°  and  150°,  a  few  short, 
thick  prisms  melting  once  at  142°  and  again  at  144°  sepa- 
rated, leaving  the  compound  which  melts  at  193°.  Of  these 
crystals  there  was  not  enough  for  a  combustion,  but  an 
analysis  of  a  fraction  melting  almost  completely  at  137°  gave 
figures  which  show  without  doubt  that  this  portion  had  the 
same  percentage  composition  as  the  body  melting  at  193°.  If 
we  are  to  regard  the  product  as  it  is  first  obtained  as  a  mixture 
of  two  ethers,  and  consider  that  142° -144°  is  the  melting-point 
of  the  lower  body,  we  can  account  for  the  melting  of  fractions 
below  142°  on  the  ground  that  mixtures  may  have  a  lower 
melting-point  than  either  of  the  component  substances. 
These  results  were  obtained  repeatedly  from  three  separate 
portions  of  Kahlbaum's  C.  P.  orthocresol,  purchased  at 
different  times,  as  well  as  from  toluquinone  melting  at  67° 
C.  and  wholly  volatile  when  exposed  to  the  air. 

When  the  silver  salt  of  toluquinonemetaoxime  acts  upon 
methyl  iodide,  the  product  is,  likewise,  not  a  single  ether 
but  a  mixture  of  ethers,  softening  at  55°  and  not  melting 
completely  below  70°.  From  it  a  body  melting  at  73° -74° 
can  easily  be  separated,  but  no  other  compound  of  definite 
melting-point  or  different  crystal  form  could  be  obtained. 

The  acetyl  compound  of  toluquinonemetaoxime  presents 
phenomena  similar  to  those  of  the  benzoyl  ethers,  but  in 
rather  more  marked  degree.  The  product,  as  first  obtained 
by  the  action  of  acetyl  chloride  on  the  silver  salt  of  the 
oxime,  or  by  acetic  anhydride  on  the  oxime  itself,  begins  to 
soften  at  90°  and  melts  completely  at  110°.  Upon  the  first 
recrystallization  a  distinction  in  crystal  form  appears,  and  a 


150  THE  ETHERS  OF  TOLUQUINONEOXIME. 

second  crystallization  of  the  separated  portions  gave  short, 
thick  prisms  melting  at  112°-113°,  and  some  smaller,  spheri- 
cally grouped  crystals,  melting  at  85° -87°.  The  extreme 
difficulty  of  preparing  the  acetyl  compound  prevented  further 
investigation. 

In  like  manner,  the  benzoyl  ether  of  monobromtoluquinone- 
metaoxime  seems  to  be  a  mixture  of  isomeric  bodies.  Thus,  it 
was  found  possible  to  add  two  atoms  of  bromine  to  the  ethers 
of  toluquinoneoxime  forming  colorless  dibromides  correspond- 
ing to  the  dibrom  addition-products  of  quinoneoxime,  and 
these  dibromides,  when  boiled  with  dilute  alcohol,  split  off 
hydrobromic  acid  with  the  formation  of  colored  monobrom 
substitution-products.  The  benzoyl  ether  of  monobromtolu- 
quinonemetaoxime  thus  formed  shows  a  variation  in  melting- 
point  similar  to  that  of  the  ethers  previously  discussed. 

The  foregoing  facts  speak  in  favor  of  Kehrmann's  theory  of 
space  isomerism  in  the  oximes  so  far  as  the  metaoximes  are 
concerned.  On  the  other  hand,  there  is  no  evidence  to  show  the 
presence  of  isomers  in  the  ethers  of  toluquinoneorthooxime : 
the  methyl,  acetyl,  and  benzoyl  ethers  all  act  as  simple 
substances,  each  product  melting  completely  at  a  definite 
temperature.  It  is  difficult  to  understand  why  isomerism 
should  be  so  much  more  evident  in  the  ethers  of  toluquinone- 
metaoxime  than  in  the  ethers  of  the  orthooximes,  unless, 
possibly,  the  closer  proximity  of  the  side-chain  to  the 
oximido-group  prevents  the  formation  of  a  space  isomer. 
There  is,  however,  a  remote  possibility  that  isomeric  bodies 
may  exist,  so  similar  in  properties  that  they  cannot  be  detected 
by  the  ordinary  methods. 

EXPERIMENTAL  PART. 
Preparation  of  the  Oximes  and  their  Salts. 

The  toluquinone-,  ortho-,  and  metaoximes  used  in  the 
experiments  to  be  described  were  made  hi  the  following 
manner :  To  a  solution  of  10  grams  of  cresol  and  8  grams  of 
potassium  nitrite  in  900  cm3  of  water,  a  solution  of  6  grams 


THE  ETHERS  OF  TOLUQUINONE  OXIME.  151 

of  concentrated  sulphuric  acid  in  100  cm3  of  water  was  added 
in  small  portions  during  the  course  of  half  an  hour,  care 
being  taken  that  both  original  solutions  should  be  between  5° 
and  10°,  and  that  this  temperature  be  maintained  during  the 
mixing.  Nearly  all  of  the  oxime  separates  out  on  standing 
in  ice-water  for  one  to  two  hours,  and  after  filtering  and 
washing  with  200-300  cm3  of  ice-water,  the  amount  of  oxime 
obtained  by  extracting  the  filtered  solution  with  ether  is  so 
small  that  it  may  be  disregarded.  The  substances  were 
purified  by  dissolving  in  a  saturated  solution  of  sodium 
carbonate  and  filtering  into  dilute  sulphuric  acid,  cooled  with 
ice.  At  this  stage  it  is  generally  ready  for  use,  but  if  further 
purification  is  desired,  it  may  be  accomplished  by  dissolving 
the  oxime  in  ether  and  shaking  with  animal  charcoal.  Upon 
filtering  and  evaporating,  the  oxime  crystallizes  in  long, 
slightly  colored  needles.  The  yield  is  large  in  both  cases, 
that  of  the  orthooxime  being  nearly  theoretical.  As  given  by 
Beilstein,  toluquinone-ra-oxime  melts  at  134°  C.  Toluquin- 
one-o-oxime  melts  at  155°  C. ;  Beilstein  gives  145°-150°. 

The  silver  salt  of  the  metaoxime  was  made  in  the  following 
manner :  5  grams  of  toluquinone-7?M)xime  were  dissolved  in 
a  solution  of  sodium  hydroxide,  a  little  less  than  the  quantity 
calculated  to  form  the  sodium  salt,  and  this  solution  was 
filtered  into  500  cm3  of  water  containing  1J  times  the  calcu- 
lated amount  of  silver  nitrate.  The  precipitate  comes  down 
in  a  flocky,  gelatinous  condition,  but  goes  over  into  a  granular 
form  on  heating  to  50°  in  a  water-bath.  When  dissolved  in 
the  least  possible  amount  of  warm  dilute  ammonia,  and  the 
solution  precipitated  with  hydrochloric  acid, 

0.1006  gram  of  the  substance,  dried  over  H2S04,  gave  0.0590 
gram  AgCl. 

Calculated  for  V^^A 

C7H6N02Ag. 

Ag  44.23  44.14 

The  salt  is  light  reddish-brown  when  first  formed,  but 
turns  darker  on  standing  or  heating.  It  decomposes  when 
heated  to  100°  and,  when  thoroughly  dry,  is  spontaneously 


152         THE  ETHERS  OF  TOLUQUINONEOXIME. 

inflammable  at  a  temperature  above  60°.     It  is  a  rather  un- 
stable body  and  cannot  be  kept  long  when  at  all  impure. 

The  silver  salt  of  toluquinone-o-oxime  was  made  in  the 
manner  described  for  the  preparation  of  the  silver  salt  of 
the  meta  form.  The  salt  falls  in  the  cold  as  reddish-brown 
crystals,  tending  to  darken  when  exposed  to  light  or  heat. 

0.2560  gram,  dried  over  H2S04,  gave  0.1498  gram  AgCl. 

Calculated  for  ,,        , 

C7H6N02Ag.  Found' 

Ag  44.23  44.06 

Although  very  similar  in  all  its  properties  to  the  silver  salt 
of  toluquinone-w-oxime,  this  salt  is  like  all  the  ethers  of  the 
ortho  form,  much  more  stable  than  its  corresponding  meta 
isomer. 

Toluquinone-m-oxime  Methyl  Ether. 

From  o-  CresoL  —  Of  the  silver  salt  of  toluquinone-w-oxime, 
2  or  3  grams  were  suspended  in  10-15  cm3  of  ligroin  and 
twice  the  calculated  quantity  of  methyl  iodide  added.  After 
standing  for  an  hour  with  frequent  shaking,  the  liquid  was 
filtered  off  and  the  residue  extracted  with  a  little  hot  ligroin. 
The  united  ligroin  solutions  were  allowed  to  evaporate  spon- 
taneously, and  the  methyl  ether  came  out  in  large,  dark- 
yellow,  hexagonal  prisms.  A  little  more  may  be  obtained  by 
allowing  the  residue  to  stand  for  a  week  with  methyl  iodide. 
The  yield  in  any  case  is  small,  the  best  results  apparently 
being  obtained  by  using  not  more  than  2  or  3  grams  of  the 
silver  salt  at  one  tune.  After  purifying  with  animal  charcoal 
and  recrystallizing  from  ligroin,  the  product  obtained  softens 
at  55°  yet  does  not  melt  completely  below  70°.  Portions 
melting  at  73°-74°  C.  were  separated  by  fractional  crystalli- 
zation, and,  on  analysis, 

0.1101  gram  of  this  body,  dried  over  H2S04,  gave  0.2582  gram 
C02  and  0.0599  gram  H20. 

0.0881  gram  of  the  substance  gave  7.1  cm8  N  at  15°  C.  and 
772  mm.  pressure. 


THE  ETHERS  OF  TOLUQUINONEOXIME.         153 

Calculated  for  v>nm<{ 

CgHoNO,. 

C  63.53  63.96 

H  6.00  6.05 

C  9.29  9.59 

Although  fractions  were  often  obtained  melting  from  55°- 
60°,  no  other  compound  of  very  definite  melting-point  could 
be  separated.  The  methyl  ether  is  very  soluble  in  all  organic 
reagents  ;  in  hot  ligroin  it  is  much  more  soluble  than  in  cold, 
from  which  it  crystallizes  in  small  bright-yellow  prisms. 

From  Toluquinone.  —  To  a  solution  of  2  grams  of  tolu- 
quinone  in  800  cm3  of  water,  the  calculated  amount  of 
methoxylamine  hydrochloride  was  added.  Yellow  crystals 
began  to  precipitate  in  the  course  of  two  hours,  and  at  the 
end  of  twelve  hours  the  reaction  was  completed.  The  liquid 
was  filtered  off  and  extracted  with  ether,  which,  upon  evapo- 
ration, left  behind  a  yellow  crystalline  mass.  The  two 
portions  were  united  and  recrystallized  from  ligroin.  The 
yield  was  very  good,  being  75  per  cent  of  the  theory.  Even 
after  boiling  in  ligroin  with  animal  charcoal  and  recrystallizing 
several  tunes,  the  substance  acts  like  a  mixture,  softening  at 
58°  and  melting  at  70°.  A  portion,  less  soluble  than  any 
other,  was  easily  separated,  which  melted  at  73°-74°  C.  and 
was  identical  in  all  respects  with  the  methyl  ether  obtained 
from  the  silver  salt  of  toluquinone-m-oxime  made  from 
0-cresol. 

0.1263  gram,  dried  over  H2S04,  gave  0.2962  gram  C02  and 
0.0672  gram  H20. 

0.2952  gram  gave  25  cm8  N  at  15°  C.  and  772  mm.  pressure. 


C  63.53  63.96 

H  6.00  5.91 

N  9.29  10.00 

Toluquinonemetaoxime  Acetyl  Ether. 

This  ether  can  be  made  in  two  ways  :  By  adding  the  cal- 
culated  amount  of  acetyl  chloride,   drop   by  drop,  to   2-3 


154  THE  ETHERS   OF  TOLUQUINONEOXIME. 

grams  of  the  silver  salt  suspended  in  15-20  cm8  of  ligroin  or 
absolute  ether,  kept  cool  by  ice-water,  evaporating  at  once, 
and  extracting  the  residue  with  hot  ligroin;  or  by  heating 
1  molecule  of  the  oxime  on  the  water-bath  for  an  hour  with 
1.5  molecules  of  acetic  anhydride,  adding  cold  water,  filtering 
off  the  tarry  mass  which  separates,  and  extracting  it  with  hot 
ligroin.  The  yield  by  either  method  is  extremely  poor,  and 
sometimes  after  purifying  by  boiling  with  animal  charcoal, 
the  total  product  consisted  of  a  few  small  crystals.  By  frac- 
tional crystallization  two  portions  were  separated,  the  less 
soluble  composed  of  thick  irregular  prisms  melting  at  112°- 
113°,  and  a  much  smaller  fraction  of  minute,  spherically 
grouped  crystals,  melting  at  85°-87°.  In  analyzing  the 
original  product  unf ractioned  : 

0.0953  gram,  dried  over  H2S04,  gave  0.2090  gram  C02  and 
0.0444  gram  H20. 

0.0848  gram  gave  5.8  cm8  N  at  15°  C.  and  760  mm.  pressure. 

Calculated  for  -    „, 

C9H9N03.  Found. 

C  60.30  59.91 

H  5.06  5.18 

N  7.84  8.01 

The  acetyl  ether  is  very  soluble  in  alcohol  and  ether,  much 
less  in  ligroin,  and  very  little  soluble  in  water. 

Toluquinone-m-oxime  Benzoyl  Ether. 

From  o-Cresol.  —  This  ether  can  be  made  from  the  silver  salt 
suspended  in  absolute  ether,  or,  better,  from  the  sodium  salt 
in  alcohol  solution.  Slightly  less  than  the  amount  of 
sodium  calculated  to  form  a  sodium  oxime  is  dissolved  in  100 
cm3  of  alcohol,  2-5  grams  of  the  oxime  added,  and  into  the 
filtered  solution  slightly  more  than  the  theoretical  quantity  of 
benzoyl  chloride  is  dropped  slowly,  the  solution  being  kept 
cool.  The  benzoyl  ether  begins  to  separate  immediately,  and 
after  a  few  moments  the  alcohol  can  be  filtered  off  and 
rejected,  as  it  contains  little  of  the  substance.  After  boiling 


THE  ETHERS  OF  TOLUQUINONEOXIME.         155 

in  alcohol  with  animal  charcoal,  when  the  solution  is  submitted 
to  fractional  crystallization,  three-fourths  of  the  crude  product 
can  be  readily  separated  in  the  form  of  bright-yellow  needles 
melting  at  193°.  Upon  concentrating  the  mother-liquor  to 
a  small  volume  and  cooling,  nearly  the  theoretical  quantity  of 
the  benzoyl  ether  can  be  recovered.  Nothing  further  was 
ever  obtained  save  a  few  flakes  of  benzoic  acid  formed  by  the 
saponifying  action  of  hydrochloric  acid,  produced  by  the 
slight  excess  of  benzoyl  chloride  acting  on  alcohol.  When 
the  portion  obtained  upon  concentration  was  repeatedly 
fractioned,  it  could  be  separated  into  portions  melting  approxi- 
mately at  193°,  and  others  melting  almost  completely  from  as 
low  as  135°  to  155°.  There  seemed  to  be  a  tendency,  however, 
for  these  lower  fractions  to  liquefy  at  142°  -144°,  and  once, 
from  an  alcoholic  solution  of  a  fraction  melting  at  140°-150°, 
that  evaporated  at  ordinary  temperature,  short,  thick,  prismatic 
crystals  separated  from  the  curved  needles  of  the  higher 
melting  (193°)  fraction.  These  few  prisms  melted  at  144° 
without  decomposition.  An  analysis  of  the  body  liquefying 
at  193°  C.  gave  the  following  figures: 

0.2277  gram,  dried  over  H2S04,  gave  0.5785  gram  C02  and 
0.0945  gram  H20. 

0.2001  gram  gave  10  cm8  N  at  15°  C.  and  760  mm.  pressure. 


C                      69.68  69.30 

H                        4.60  4.61 

N                       5.82  5.85 

An  analysis  of  a  fraction   melting  from  145°  -165°  gave 
these  percentages: 

0.1421  gram,  dried  over  H2S04,  gave  0.3635  gram  C02  and 
0.0565  gram  H20. 

0.6726  gram  gave  34.1  cm8  N  at  15°  C.  and  770  mm.  pressure. 

Calculated  for  uv*,,^ 

n   TT  icrk  iround. 

C14HnNO3. 

C                       69.68  69.74 

H                        4.60  4.43 

N                        5.82  6.02 


156         THE  ETHERS  OF  TOLUQUINONEOXIME. 

The  benzoyl  ether  is  not  at  all  soluble  in  water  or  ligroin,  is 
slightly  soluble  in  cold  alcohol,  but  dissolves  readily  in  ether, 
chloroform,  and  glacial  acetic  acid.  The  low-melting  body 
seems  to  be  more  soluble  in  alcohol  than  its  higher-melting 
isomer,  but,  as  nine-tenths  of  the  total  product  appeared  to 
be  the  body  melting  at  193°,  the  small  proportion  of  the  low- 
melting  compound  may  account  partially  for  the  idea  that  the 
latter  is  more  soluble. 

From  Toluquinone.  —  The  oxime  was  made  according  to  the 
method  of  Goldschmidt  and  Schmidt,*  by  treating  toluquinone 
in  aqueous  solution  with  an  excess  of  hydroxylamine  hydro- 
chloride,  extracting  with  ether,  and  purifying  with  animal 
charcoal.  From  the  sodium  salt  of  the  oxime  thus  formed 
the  benzoyl  ether  was  made  in  the  manner  previously  described. 
The  product  is  identical  with  the  benzoyl  ether  made  from 
o-cresol.  The  body  melting  at  193°  C.  was  readily  isolated 
and  analyzed: 

0.1002  gram,  dried  over  H2S04,  gave  0.2565  gram  C02  and 
0.0452  gram  H20. 

0.3581  gram  gave  17.8  cm8  N  at  15°  C.  and  760  mm.  pressure. 

Calculated  for  v       , 

CMHUN08.  Found. 

C  69.68  69.81 

H  4.60  5.01 

N  5.82  5.77 

By  fractional  crystallization,  low-melting  portions  were 
separated,  exactly  as  hi  the  case  of  the  benzoyl  ether  made 
from  o-cresol,  except  that  some  fractions  were  obtained  melt- 
ing  partially  as  low  as  129°.  In  order  to  be  certain  that 
benzoic  acid  (m.  p.,  120°  C.)  was  not  unduly  lowering  these 
melting-points,  these  fractions  were  boiled  with  water  and 
filtered  hot.  The  filtrate  was  not  acid  to  litmus  and  contained 
only  a  trace  of  organic  matter.  The  ethers,  when  dried,  gave 
the  same  melting-point  as  before  boiling  with  water,  and, 
upon  recrystallizing  from  a  little  alcohol,  did  not  exhibit  any 

*  Ber.  Dtsch.  chem.  Ges.,  xvii,  2063. 


THE  ETHERS  OF  TOLUQUINONEOXIME.         157 

change.  Furthermore,  the  benzoyl  ether  obtained  by  the 
usual  method,  when  only  75  per  cent,  of  the  theoretical 
quantity  of  benzoyl  chloride  was  used,  gave  fractions  begin- 
ning to  melt  at  129°.  Obviously  benzoic  acid  could  not  be 
present  in  these  instances. 

By  slow,  spontaneous  evaporation  of  the  alcoholic  solution 
of  a  fraction  liquefying  at  140°-150°,  a  few  crystals  melting 
at  142°,  similar  to  those  obtained  in  the  same  manner  from 
the  oxime  made  from  0-cresol,  separated  from  the  body 
melting  at  193°.  Since  there  was  not  enough  of  this  com- 
pound for  a  combustion,  an  analysis  was  made  of  a  fraction 
melting  at  137°  C.  with  the  following  results : 

0.1131  gram,  dried  over  HgSO^  gave  0.2901  gram  C02  and 
0.0480  gram  H20. 

Calculated  for  _       , 

CUHUN08.  Found- 

C  69.68  69.96 

H  4.60  4.72 

Fractions  were  frequently  obtained  melting  at  about  177° 
C.,  and  from  these  attempts  were  repeatedly  made  to  isolate 
some  of  the  ortho  isomer.  None  was  ever  detected  under  the 
lens,  and  fractional  crystallization  always  separated  such 
portions,  principally  into  the  body  melting  at  190°,  and  a 
small  fraction  melting  much  lower. 

Since  the  benzoyl  ether  will  decompose  into  a  dark-brown 
liquid  with  the  evolution  of  brown  fumes  of  nitrogen  oxides 
when  heated  above  160°,  it  is  only  by  rapidly  heating  that 
193°  can  be  observed  as  a  melting-point.  This  applies  also 
to  the  bromine  addition-products  of  the  benzoates  of  both 
the  ortho-  and  metaoximes. 

Dibromtoluquinone-m-oxime  Benzoyl  Ether. 

The  benzoyl  ether  was  dissolved  in  chloroform  and  cooled 
while  the  theoretical  quantity  of  bromine  was  added  in  small 
portions.  After  standing  for  an  hour,  the  chloroform  was 
evaporated  spontaneously,  and  the  light-brown  residue  re- 


158  THE  ETHERS  OF  TOLUQUINONEOXIME. 

crystallized  from  glacial  acetic  acid.  It  can  also  be  purified 
by  dissolving  in  fuming  nitric  acid  and  pouring  into  water. 
On  analysis : 

0.1437  gram,  dried  over  H2S04,  gave  0.2176  gram  C02  and 
0.0339  gram  H20. 

0.1215  gram  gave  0.1140  gram  AgBr. 

Calculated  for  ,,       , 

CMHuBr2N08.  Found- 

C  41.90  41.30 

H  2.76  2.62 

Br  39.87  39.93 

The  dibromide  is  insoluble  in  water,  somewhat  soluble  in 
cold  alcohol,  and  readily  dissolves  in  chloroform  and  glacial 
acetic  acid,  from  which  it  crystallizes  in  white  prisms  melting 
at  165°  C.  with  decomposition. 

Monobromtoluquinone-m-oxime  Benzoyl  Ether. 

When  the  dibromide  is  boiled  with  alcohol  hydrobromic 
acid  splits  off,  two  hours  being  required  to  complete  the 
process,  during  which  little  or  no  saponification  takes  place. 
After  recrystallizing  from  alcohol,  a  mixture  of  monobrom 
compounds  is  obtained,  melting  from  155° -170°.  A  portion 
melting  with  decomposition  at  174°  C.,  was  separated  and 
analyzed  : 

0.1208  gram,  dried  over  H2S04,  gave  0.2319  gram  C02  and 
0.0345  gram  H20. 

0.0735  gram  gave  0.0425  gram  AgBr. 

Calculated  for  Fminrl 

C14H10BrN08. 

C  52.49  52.36 

H  3.15  3.17 

Br  24.98  24.63 

It  is  very  similar  in  its  properties  to  toluquinone-ra-oxime 
benzoyl  ether,  crystallizing  from  alcohol  in  bright>yellow 
needles. 


THE  ETHERS  OF  TOLUQUINONEOXIME.  159 

Toluquinone-o-oxime  Methyl  Ether. 

The  methyl  ether  of  toluquinone-o-oxime  was  made  from 
the  silver  salt  and  methyl  iodide  in  the  same  manner  as  its 
meta  isomer,  the  yield  being  somewhat  better.  Once  recrys- 
tallized  from  ligroin: 

0.1247  gram  of  the  ether,  dried  over  H2S04,  gave  0.2931  gram 
C02  and  0.0675  gram  H20. 

0.3484  gram  gave  27  cm8  N  at  15°  C.  and  772  mm.  pressure. 

Calculated  for  •„,       •, 

Found' 


C  63.53  64.10 

H  6.00  6.01 

N  9.29  9.22 

Its  properties  are  almost  identical  with  the  methyl  ether 
of  the  meta  form.  It  crystallizes  from  ligroin  in  long  yellow 
needles,  every  portion  of  which,  obtained  by  fractional  crys- 
tallization, melts  at  69°  C. 

Dibromtoluquinone-o-oxime  Methyl  Ether. 

The  methyl  ether  was  dissolved  in  chloroform,  cooled,  and 
the  calculated  quantity  of  bromine  added.  The  reaction  is 
completed  in  twenty  minutes,  and,  on  evaporation  of  the 
chloroform,  the  dibromide  is  left  behind  as  a  dirty  white 
mass.  Once  recrystallized  from  ligroin,  the  substance  is 
ready  for  analysis: 

0.1690  gram,  dried  over  H2S04,  gave  0.1877  gram  C02  and 
0.0456  gram  H2O. 

0.0744  grain  gave  0.0909  gram  AgBr. 

Calculated  for  •&*,•,•** 

C8H9Br2N02. 

C  30.87  30.29 

H  2.92  3.00 

Br  51.41  52.00 

It  is  insoluble  hi  water  but  quite  soluble  in  most  organic 
reagents.  From  ligroin  it  crystallizes  in  white  prisms,  melt- 
ing at  112°  C. 


160  THE  ETHERS  OF  TOLUQUINONEOXIME. 

Toluquinone-o-oxime  Acetyl  Ether. 

This  ether  was  made  by  Wurster  and  Riedel  *  from  tolu- 
quinone-0-oxime  and  acetic  anhydride.  It  can  be  prepared 
also  from  the  silver  salt  and  acetyl  chloride,  the  yield  being 
very  poor,  only  a  trifle  better  than  its  meta  isomer,  which 
it  closely  resembles  in  its  properties. 

0.1194  gram,  dried  over  H2SO4,  gave  0.2666  gram  C02  and 
0.0545  gram  H2O. 

0.2389  gram  gave  15  cm8  N.  at  15°  C.  and  760  mm.  pressure. 

Calculated  for  -       , 

C9H9N08.  Found- 

C  60.30  60.89 

H  5.06  5.07 

N  7.84  7.35 

From  ligroin  it  crystallizes  in  irregular  yellow  prisms, 
melting  at  92°  C. 

Toluquinone-o-oxime  Benzoyl  Ether. 

This  ether  was  made  in  alcohol  solution  from  the  sodium 
salt  in  the  manner  already  described  for  the  benzoyl  ether  of 
the  meta  form.  The  yield  is  practically  theoretical,  and, 
after  one  recrystallization  from  alcohol,  the  product  is  ready 
for  analysis. 

0.1589  gram,  dried  over  H2S04,  gave  0.4074  gram  C02  and 
0.0686  gram  H20. 

0.5101  gram  gave  26.1  cm8  N  at  15°  C.  and  7.66  mm.  pressure. 

Calculated  for  im™,! 

CUHUN03. 

C  69.68  69.92 

H  4.60  4.80 

N  5.82  6.04 

Although  submitted  to  the  most  careful  fractional  crystal- 
lization, every  particle  obtained  melted  sharply  at  177°  C., 
with  slight  decomposition.  From  alcohol  it  crystallizes  in 
light,  brownish-yellow  blades,  which  have  all  the  properties 

*  Ber.  Dtsch.  chem.  Ges.,  xii,  1799. 


THE  ETHERS  OF  TOLUQUINONEOXIME.  161 

of  the  metaoxime  benzoyl  ether  except  that  it  is  a  little  more 
soluble  in  organic  reagents. 

Dibromtoluquinone-o-oxime  Benzoyl  Ether. 

The  benzoyl  ether  was  dissolved  in  chloroform,  bromine 
added  to  the  cooled  solution,  and  the  product  recrystallized 
from  glacial  acetic  acid. 

0.2111  gram,  dried  over  H2S04,  gave  0.3208  gram  C02  and 
0.0564  gram  H20. 

0.1003  gram  gave  0.0924  gram  AgBr. 

Calculated  for  Found 

C  41.90  41.44 

H  2.76  2.97 

Br  39.87  39.21 

It  is  insoluble  in  ligroin  and  water  and  but  little  soluble  in 
cold  alcohol.  Fuming  nitric  acid  dissolves  it  readily,  and 
from  the  solution  it  is  precipitated  unchanged  by  water.  It 
crystallizes  from  glacial  acetic  acid  hi  short,  thick,  ortho- 
rhombic  prisms  melting  at  159°  C.  with  decomposition. 


VOL.   II.  —  11 


XXI 

THE  APPLICATION  OF  IODINE  IN  THE 
ANALYSIS   OF  ALKALIES  AND  ACIDS. 

BY  CLAUDE  F.   WALKER  AND  DAVID  H.  M.  GILLESPIE* 

IT  is  well  known  that  when  a  free  mineral  acid  is  added  to  a 
neutral  mixture  of  metallic  iodate  and  iodide,  the  iodate  is 
reduced  and  iodine  is  liberated  according  to  the  equation : 

EI08  +  5EI  +  3H2S04  =  31,  +  3E2S04  +  3H20. 

This  reaction  is  complete  and  non-reversible  under  the 
conditions  of  analysis,  and  it  may  therefore  be  applied  to  the 
estimation  of  amounts  of  iodate,  iodide  or  mineral  acid  present 
in  an  unknown  solution.  A  solution  of  iodate  to  be  analyzed 
is  mixed  with  an  excess  of  iodide  and  mineral  acid,  the 
resulting  free  iodine  estimated  by  directly  titrating  with 
sodium  thiosulphate  or  arsenious  acid,  and  one-sixth  of  the 
amount  found  taken  as  equivalent  to  the  iodate  originally 
present.f  Similarly,  a  solution  of  iodide  to  be  analyzed  is 
mixed  with  an  excess  of  iodate  and  mineral  acid,  the  resulting 
free  iodine  estimated  by  directly  titrating  in  alkaline  solution 
with  arsenious  acid,  and  five-sixths  of  its  amount  taken  as 
equivalent  to  the  iodide  originally  present.^  A  solution  of 
mineral  acid  to  be  analyzed  is  mixed  with  an  excess  of  iodate 
and  iodide,  the  resulting  free  iodine  estimated  by  directly 
titrating  with  sodium  thiosulphate,  and  its  entire  amount 
taken  as  equivalent  to  the  amount  of  mineral  acid  originally 
present.  §  Groger  has  applied  the  last  mentioned  method  to 

*  From  Am.  Jour.  Sci.,  vi,  455. 

t  Rammelsberg,  Fogg.  Ann.,  cxxxv,  493 ;  Walker,  Am.  Jour.  Sci.,  iv,  235. 
This  volume,  p.  52. 

J  Gooch  and  Walker,  Am.  Jour.  Sci.,  iii,  293.    This  volume,  p.  33. 

§  Kjeldahl,  Zeitschr.  anal.  Chem.,  xxii,  366 ;  Furry,  Am.  Chem.  Jour.,  vi, 
341 ;  Groger,  Zeitschr.  angew.  Chem.,  1894,  52. 


IODINE  IN  ANALYSIS  OF  ALKALIES,  ETC.         163 

the  direct  analysis  of  various  mineral  acids,  and  has  obtained 
results  manifestly  better  than  those  afforded  by  the  use  of 
vegetable  indicators.  Groger  has  also  indirectly  analyzed 
solutions  of  alkali  hydroxides  and  carbonates  by  adding  the 
solution  to  be  analyzed  to  a  measured  volume  of  mineral  acid, 
previously  standardized  by  the  above  method,  and  estimating 
the  small  excess  of  free  mineral  acid  that  finally  remains  by 
the  same  method.  The  only  difficulty  with  the  Groger 
process  lies  in  the  fact  that  under  the  conditions  present  the 
end-point  of  the  final  reaction  between  iodine  and  sodium 
thiosulphate  is  somewhat  obscured  by  a  peculiar  back-play 
of  color  due  to  a  continuous  slow  liberation  of  iodine  in  the 
system. 

When  a  solution  of  a  metallic  hydroxide  is  acted  on  by 
iodine  at  a  temperature  high  enough  to  decompose  the  small 
amounts  of  hypoiodites  that  might  otherwise  be  present, 
the  final  action  results  in  the  formation  of  an  exactly  neutral 
mixture  of  iodate  and  iodide,  according  to  the  equation : 

6EOH  +  3I2  =  BI08  +  5RI  +  3H20. 

Phelps  *  has  shown  that  in  the  case  of  barium  hydroxide  at 
least  this  reaction  is  regular  and  complete  under  the  conditions 
of  analysis,  and  is  independent  of  the  excess  of  iodine  which 
remains  hi  the  neutral  mixture  unacted  upon  and  may  be 
estimated  by  directly  titrating  with  arsenious  acid.  Phelps 
not  only  applies  this  principle  of  action  to  the  standardization 
of  solutions  of  barium  hydroxide  by  boiling  with  an  excess  of 
iodine  in  a  trapped  flask,  but  also  bases  thereon  a  differential 
method  for  determining  carbon-dioxide,  in  which  the  liberated 
gas  is  run  into  a  measured  amount  of  barium  hydroxide,  the 
final  excess  of  which  is  estimated  by  treating  with  iodine  in 
the  presence  of  the  precipitated  barium  carbonate.  The  good 
result  obtained  by  Phelps  with  barium  hydroxide  suggested 
that  the  attempt  be  made  to  analyze  alkali  hydroxides,  and 
possibly  carbonates,  by  a  method,  simpler  than  that  devised  by 
Groger,  based  on  the  direct  treatment  of  these  compounds 

*  Am.  Jour.  Sci.,  ii,  70.    Volume  I,  p.  369. 


164  APPLICATION  OF  IODINE  IN  THE 

with  iodine  in  hot  solution.  It  also  seemed  possible  to  apply 
the  differential  method  not  only  to  carbon  dioxide  but  to  any 
acid  or  other  compound  that  will  act  definitely  and  completely 
with  the  metallic  hydroxide  employed,  provided  the  soluble  or 
insoluble  product  formed  will  not  be  attacked  when  heated  in 
the  presence  of  iodine.  It  was  decided  to  modify  the  Phelps 
process,  however,  in  order  to  obviate  the  necessity  of  handling 
large  measured  amounts  of  iodine  in  a  flask  trapped  to  prevent 
mechanical  loss  by  heating.  The  flask  was  therefore  dispensed 
with  altogether,  and  the  hydroxide  solution  to  be  analyzed 
was  mixed  with  an  approximately  measured  excess  of  iodine 
solution,  in  an  Erlenmeyer  beaker,  the  mouth  of  which  was 
lightly  closed  with  a  little  trap  to  prevent  loss  by  spattering. 
The  excess  of  iodine  was  then  completely  removed  by  boiling, 
and  the  cooled  colorless  solution  remaining,  which  contained 
a  neutral  mixture  of  iodate  and  iodide,  was  acidified  with  a 
mineral  acid  and  the  liberated  iodine  titrated  with  sodium 
thiosulphate,  the  amount  found  being  equivalent  to  the 
amount  of  hydroxide  taken  for  analysis. 

The  present  investigation  was  undertaken  to  study  the 
limitations  and  possible  applications  in  analysis  of  the 
reactions  between  iodine  on  the  one  hand,  and  barium 
hydroxide,  potassium  hydroxide  and  sodium  carbonate  on  the 
other.  It  was  soon  found  that  the  reaction  in  the  case  of 
sodium  carbonate  is  entirely  dependent  on  conditions  of  time, 
mass,  and  temperature,  and  cannot  be  pushed  to  completion 
except  under  conditions  that  make  its  application  in  analysis 
impossible.  In  the  case  of  barium  and  potassium  hydroxides 
both  the  original  procedure  of  Phelps  and  the  modification 
above  described  were  employed.  The  modified  method  was 
found  to  be  the  more  convenient  and  speedy  of  the  two.  The 
results  obtained  in  both  cases  agreed  with  one  another, 
but  were  invariably  lower  by  a  small  nearly  constant 
amount  than  those  obtained  by  both  the  gravimetric  and  the 
Groger  processes.  This  error  of  the  Phelps  process  and  its 
modification  is  possibly  due  to  the  action  of  atmospheric  carbon 
dioxide  on  the  hydroxide  solution  during  the  short  tune  it  is 


ANALYSIS  OF  ALKALIES  AND  ACIDS.  165 

exposed.  While  it  will  affect  the  value  of  the  method  as  a  means 
of  accurately  determining  the  absolute  amount  of  hydroxide 
present  in  a  given  volume  of  solution,  it  cannot  so  affect  the 
accuracy  of  any  differential  method  founded  on  the  original 
Phelps  process  or  its  modification.  This  is  demonstrated  by 
the  work  of  Phelps  in  the  case  of  carbon  dioxide,  and  by  the 
present  investigation  in  the  case  of  hydrochloric  and  sulphuric 
acid.  Analyses  of  these  two  acids  were  made  by  adding  the 
solution  to  be  analyzed  to  a  measured  volume  of  barium  or 
potassium  hydroxide,  previously  standardized  by  the  modified 
Phelps  method.  The  small  excess  of  hydroxide  remaining 
was  then  estimated  by  the  same  method,  the  results  agreeing 
with  those  already  obtained  by  both  the  gravimetric  and  the 
Gro'ger  processes.  It  seems  probable  that  other  acids  and 
compounds  for  which  there  is  now  no  rapid  iodometric  method 
may  be  analyzed  by  a  method  similar  to  this,  which  has  given 
good  results  with  carbonic,  hydrochloric,  and  sulphuric  acids. 

Decinormal  solutions  of  the  alkali  hydroxides  were  prepared, 
and  kept  with  great  care  in  trapped  bottles,  from  which  por- 
tions for  analysis  were  measured  by  means  of  a  self-feeding 
burette,  which  was  also  fitted  with  a  trap.  All  vessels  and 
water  used  were  made  as  free  as  possible  from  carbon  dioxide, 
and  the  operations  were  conducted  as  rapidly  as  possible. 

In  the  analyses  by  the  Phelps  method  a  carefully  measured 
excess  of  decinormal  iodine  was  drawn  into  a  small  ether  wash- 
bottle,  and  the  desired  amount  of  alkali  was  rapidly  run  into 
it.  The  stopper,  to  which  had  been  sealed  a  Will  and  Varren- 
trapp  absorption  bulb  was  placed  in  the  bottle  and  the  bulb 
was  charged  with  a  5  per  cent  solution  of  potassium  iodide  to 
catch  any  escaping  vapors  of  iodine.  The  apparatus  was 
placed  over  a  low  flame  and  the  contents  heated  to  boiling  or 
slightly  longer,  and  then  cooled  hi  a  stream  of  water.  The 
contents  of  the  bulb  and  connecting  tubes  were  then  washed 
into  the  flask,  and  the  excess  of  free  iodine  remaining  was 
titrated  with  arsenious  acid,  in  the  presence  of  5  cm3  of  starch 
emulsion.  Blank  analyses  were  made  to  insure  against  me- 
chanical loss  of  iodine  during  boiling  and  to  prevent  any  error 


166 


APPLICATION  OF  IODINE  IN  THE 


on  account  of  the  presence  of  carbonate  or  other  impurity  in 
the  solutions  employed. 

Some  of  the  results  obtained  with  barium  hydroxide  are 
given  in  Table  I.  The  variation  in  different  analyses  of  the 
same  series  is  not  large,  and  the  results  are  independent  of 
the  amount  taken  for  analysis  and  of  the  excess  of  iodine 

employed. 

TABLE  I. 

ANALYSES  OP  *0  BARIUM  HYDROXIDE  SOLUTION. 
(By  boiling  in  a  trapped  flask  with  an  excess  of  iodine.) 


Exp. 

Ba(OH)a 
taken. 

Iodine 
taken. 

Iodine 
absorbed  by 

Ba(OH)3 
found. 

Mean. 

Variation. 

Ba(OH)2. 

cm8 

grm. 

grm. 

grm. 

grm. 

grm 

(1) 

10 

0.13 

0.1054 

0.0712 

0.0699 

0.0013+ 

(2) 

10 

0.14 

0.1028 

0.0692 

0.0699 

0.0007- 

(3) 

20 

0.23 

0.2072 

0.1399 

0.1398 

0.0001+ 

(4) 

20 

0.25 

0.2074 

0.1401 

0.1398 

0.0003+ 

(5) 

40 

0.44 

0.4143 

0.2798 

0.2796 

0.0002+ 

(6) 

40 

0.44 

0.4148 

0.2802 

0.2796 

0.0006+ 

(7) 

40 

0.48 

0.4160 

0.2809 

0.2796 

0.0013+ 

(8) 

40 

0.48 

0.4126 

0.2786 

0.2796 

0.0010- 

(9) 

40 

0.51 

0.4115 

0.2779 

0.2796 

0.0017- 

(10) 

40 

0.51 

0.4136 

0.2793 

0.2796 

0.0003- 

The  analyses  of  potassium  hydroxide  were  made  in  the 
same  way  as  were  those  of  barium  hydroxide,  and  gave  quite 
similar  results.  They  follow  in  Table  II. 

TABLE  H. 

ANALYSIS  OP  J  POTASSIUM  HYDROXIDE  SOLUTION. 
(By  boiling  in  a  trapped  flask  with  an  excess  of  iodine.) 


Exp. 

KOH 

taken. 

Iodine 
taken. 

Iodine 
absorbed  by 
KOH. 

KOH 

found. 

Mean. 

Variation. 

cm> 

grm. 

grm. 

grin* 

grm. 

grm. 

(1) 

10 

0.20 

0.1621 

0.0716 

0.0717 

0.0001- 

(2) 

10 

0.23 

0.1613 

0.0715 

0.0717 

0.0002- 

(3) 

15 

0.30 

0.2404 

0.1063 

0.1076 

0.0013- 

4) 

15 

0.30 

0.2429 

0.1074 

0.1076 

0.0002- 

5 

16 

0.34 

0.2431 

0.1076 

0.1076 

0.0001- 

(6) 

25 

0.51 

0.4089 

0.1808 

0.1792 

0.0016+ 

(7) 

25 

0.51 

0.4058 

0.1794 

0.1792 

O.OOOSH- 

ANALYSIS   OF  ALKALIES  AND  ACIDS.  167 

The  analyses  by  the  modification  of  the  Phelps  method  were 
made  by  drawing  into  an  Erlenmeyer  beaker  of  convenient 
size  an  approximately  measured  excess  of  decinormal  iodine, 
and  rapidly  running  the  desired  amount  of  alkali  into  it.  The 
neck  of  the  beaker  was  then  closed  by  a  little  trap,  made  of 
one  of  the  halves  of  a  double  end  calcium  chloride  drying 
tube,  to  prevent  appreciable  loss  by  spattering.  The  beaker 
was  then  placed  over  a  low  flame,  and  the  contents  boiled  until 
the  last  trace  of  the  excess  of  iodine  had  volatilized  from  the 
solution  and  the  trap.  The  volume  was  carefully  regulated 
before  and  during  the  boiling,  being  kept  as  small  as  possible, 
usually  amounting  to  about  100  cm3  at  the  start  and  35  cm3 
at  the  close.  In  the  case  of  barium  hydroxide  care  had  to  be 
taken  to  keep  the  dilution  sufficient  to  prevent  the  separation 
the  crystalline  barium  iodate,  which  is  soluble  with  difficulty. 
To  steady  the  ebullition  a  little  spiral  of  platinum  was  intro- 
duced into  the  beaker.  After  the  boiling  had  ceased,  the 
colorless  solution,  containing  a  neutral  mixture  of  iodate  and 
iodide,  was  cooled  hi  running  water  and  treated  with  10  cm3 
of  dilute  (1 : 3)  hydrochloric  acid  or  (1 : 3)  sulphuric  acid. 
The  liberated  iodine  was  titrated  directly  with  sodium  thiosul- 
phate,  hi  the  presence  of  5  cm3  of  starch  emulsion.  In  the 
case  of  barium  hydroxide  the  iodine  was  liberated  with  dilute 
(1 : 3)  hydrochloric  acid  to  save  the  inconvenience  of  working 
in  the  presence  of  precipitated  barium  sulphate ;  with  potas- 
sium hydroxide,  however,  dilute  (1 :  3)  sulphuric  acid  was 
employed.  In  view  of  a  statement  by  Pickering*  that  titra- 
tions  with  sodium  thiosulphate  in  the  presence  of  acid  involve 
an  error,  a  series  of  blank  analyses  was  made  which  showed  con- 
clusively that  no  such  error  exists  under  the  conditions  which 
obtain  in  the  process  under  consideration.  Care  was  also  taken, 
as  in  a  former  case,  to  guard  against  the  possible  presence  of 
carbonates  or  other  impurities  in  the  reagents  employed. 

In  Table  III  are  given  the  results  of  a  series  of  analyses  of 
barium  hydroxide  by  the  modified  method  just  described. 
They  agree  fairly  well  with  those  of  Table  I. 
*  Jour.  Chem.  Soc.,  xxxvii,  134. 


168 


APPLICATION  OF  IODINE  IN  THE 


ANALYSES  OF 


TABLE  III. 
BAKIUM  HYDROXIDE  SOLUTION. 


(By  boiling  with  excess  of  iodine  in  an  open  beaker  to  decoloration,  and 
acidifying  the  residue.) 


Exp. 

Ba(OH), 
taken. 

Iodine 
taken. 

Iodine 
absorbed  by 
Ba(OH),. 

Ba(OH), 
found. 

Mean. 

Variation. 

cm** 

grni. 

grin. 

gnu. 

griii. 

grin. 

(1) 

10 

0.13 

0.1023 

0.0691 

0.0695 

0.0004- 

2) 

10 

0.16 

0.1020 

0.0689 

0.0695 

0.0006- 

3 

16 

0.18 

0.1548 

0.1046 

0.1043 

0.0003+ 

4 

15 

0.20 

0.1546 

0.1045 

0.1043 

0.0002+ 

(5) 

20 

0.23 

0.2049 

0.1384 

0.1390 

0.0006- 

(6) 

20 

0.25 

0.2058 

0.1390 

0.1390 

0.0000 

(7) 

20 

0.32 

0.2065 

0.1394 

0.1390 

0.0004+ 

(8) 

25 

0.29 

0.2567 

0.1734 

0.1738 

0.0004- 

(9) 

25 

0.32 

0.2562 

0.1730 

0.1738 

0.0008- 

(10) 

40 

0.47 

0.4120 

0.2783 

0.2780 

0.0003+ 

(11) 

40 

0.48 

0.4119 

0.2782 

0.2780 

0.0002+ 

(12) 

40 

0.48 

0.4152 

0.2804 

0.2780 

0.0024+ 

(13) 

40 

0.49 

0.4109 

0.2775 

0.2780 

0.0005- 

The  analyses  of  potassium  hydroxide  by  the  modified  method 
are  given  in  Table  IV,  and  are  found  to  agree  well  with  those 
of  Table  II. 

TABLE  IV. 
ANALYSES  OF  ^  POTASSIUM  HYDROXIDE  SOLUTION. 

(By  boiling  with  excess  of  iodine  in  an  open  beaker  to  decoloration,  and 
acidifying  the  residue.) 


Eip. 

KOH 
taken. 

Iodine 
taken. 

Iodine 
absorbed  by 
KOH. 

Ba(OH)2 
found. 

Mean. 

Variation. 

cm» 

grm. 

gnu. 

grm. 

grm. 

grm. 

(1) 

10 

0.20 

0.1624 

0.0718 

0.0721 

0.0003- 

(2) 

10 

0.23 

0.1618 

0.0715 

0.0721 

0.0006- 

(3) 

10 

0.25 

0.1622 

0.0717 

0.0721 

0.0004- 

(4) 

15 

0.30 

0.2459 

0.1087 

0.1082 

0.0005+ 

(5 

15 

0.34 

0.2473 

0.1093 

0.1082 

0.0011+ 

(6) 

15 

0.38 

0.2441 

0.1079 

0.1082 

0.0003- 

(7) 

20 

0.41 

0.3274 

0.1447 

0.1442 

0.0005+ 

(8) 

20 

0.46 

0.3259 

0.1441 

0.1442 

0.0001- 

(9) 

20 

0.51 

0.3269 

0.1445 

0.1442 

0.0003+ 

(10 

25 

0.51 

0.4052 

0.1791 

0.1803 

0.0012- 

(11) 

25 

0.57 

0.4082 

0.1805 

0.1803 

0.0002+ 

(12) 

25 

0.63 

0.4080 

0.1804 

0.1803 

0.0001+ 

ANALYSIS  OF  ALKALIES  AND  ACIDS. 


169 


A  gravimetric  analysis  of  the  barium  hydroxide  solution  in 
which  the  barium  was  weighed  as  the  sulphate,  gave  as  a 
result  0.1411  grm.  Ba(OH)2  for  each  20  cm3  taken.  An  anal- 
ysis of  the  same  solution  by  the  Groger  process  gave  for  the 
same  volume  0.1420  grm.  The  result  by  the  Phelps  process, 
however,  was  0.1398  grm.,  and  by  the  modified  process  0.1390 
grm.  That  the  difference  of  2  mg.  between  the  results  by  the 
gravimetric  and  the  Groger  processes  on  one  hand,  and  the 
Phelps  process  and  its  modification  on  the  other,  may  be  due 
to  atmospheric  carbon  dioxide,  has  already  been  pointed  out. 
A  gravimetric  analysis  of  the  potassium  hydroxide  solution 
by  evaporating  and  weighing  as  KC1  gave  0.1111  grm.  KOH 
for  each  20  cm3  taken,  agreeing  with  0.1106  grm.  obtained  by 
the  Groger  process.  The  analyses  by  the  Phelps  process  and 
its  modification  for  the  same  solution  gave  0.1076  grm.  and 
0.1082  grm.  respectively.  These  results  are  strikingly  in 
accord  with  those  obtained  with  barium  hydroxide. 

In  the  application  of  the  modification  of  the  Phelps  process 
to  the  indirect  analysis  of  hydrochloric  and  sulphuric  acids  the 
procedure  was  essentially  the  same  as  that  detailed  for  the 
analysis  of  barium  and  potassium  hydroxides  in  Tables  III  and 
IV.  The  acid  solution  to  be  analyzed  was  drawn  into  an 
Erlenmeyer  beaker,  a  measured  excess  of  standardized  alkali 

TABLE  V. 
ANALYSES  OF  ^  HYDKOCHLOKIC  ACID  SOLUTION. 

(By  adding  to  excess  of  ^  Ba(OH)2,  boiling  with  excess  of  iodine  to  decolora- 
tion and  acidifying  the  residue.) 


Exp. 

HCl 

taken. 

Ba(OH), 
taken. 

Ba(OH)2 
neutralized 

HCl  found. 

Mean. 

Variation. 

by  HCl. 

cm3 

grill. 

grm. 

grm. 

grm. 

grm. 

(1) 

15 

0.17 

0.1128 

0.0480 

0.0476 

0.0004+ 

2 

15 

0.17 

0.1118 

0.0475 

0.0476 

0.0001- 

(3) 

16 

0.17 

0.1112 

0.0473 

0.0476 

0.0003- 

(4) 

25 

0.26 

0.1860 

0.0791 

0.0794 

0.0003- 

5) 

25 

0.26 

0.1866 

0.0794 

0.0794 

0.0000 

6 

35 

0.34 

0.2634 

0.1120 

0.1111 

0.0009+ 

(7) 

36 

0.34 

0.2603 

0.1107 

0.1111 

0.0004- 

170 


APPLICATION  OF  IODINE  IN  THE 


added,  and  the  operation  completed  as  described.  It  was 
found  that  barium  hydroxide  and  potassium  hydroxide  may  be 
applied  with  equal  accuracy  to  the  analysis  of  both  hydrochloric 
and  sulphuric  acids.  Some  of  the  results  obtained  are  given 
in  Tables  V,  VI  and  VII. 

TABLE  VI. 

ANALYSES  OP  ^  HYDROCHLORIC  ACID  SOLUTION. 

(By  adding  to  excess  of  ^  KOH,  boiling  with  excess  of  iodine  to  decoloration, 
and  acidifying  the  residue.) 


Exp. 

HC1 

taken. 

KOH 

taken. 

KOH 

neutralized 
by  HC1. 

HC1  found. 

Mean. 

Variation. 

cm8. 

grm. 

grm. 

grm. 

grm. 

grm. 

(1) 

20 

0.14 

0.0972 

0.0633 

0.0633 

0.0000 

(2) 

20 

0.14 

0.0975 

0.0634 

0.0633 

0.0001+ 

(3) 

25 

0.14 

0.1222 

0.0795 

0.0791 

0.0004+ 

(4) 

25 

0.14 

0.1207 

0.0785 

0.0791 

0.0006- 

TABLE  VII. 

ANALYSES  OF  J  SULPHURIC  ACID  SOLUTION. 

(By  adding  to  excess  of  ^  Ba(OH)2,  boiling  with  excess  of  iodine  to  decolora- 
tion, and  acidifying  the  residue.) 


Eip. 

as 

BaOH2 
taken. 

Ba(OH), 
neutralized 
by  H2804. 

H2804 
found. 

Mean. 

Variation. 

cm8 

grm. 

gTIH. 

grm. 

grm. 

grm. 

(1) 

10 

0.21 

0.0884 

0.0506 

0.0498 

0.0008+ 

(2) 

10 

0.21 

0.0880 

0.0503 

0.0498 

0.0005+ 

(3) 

15 

0.30 

0.1328 

0.0754 

0.0748 

0.0006+ 

(4) 

15 

0.30 

0.1313 

0.0751 

0.0748 

0.0003+ 

(5) 

25 

0.43 

0.2168 

0.1239 

0.1246 

0.0007- 

(6) 

30 

0.43 

0.2600 

0.1481 

0.1495 

0.0014- 

An  analysis  of  the  hydrochloric  acid  solution  by  the  Grb'ger 
method,  which  was  found  to  agree  in  every  case  with  the 
gravimetric  determination,  gave  for  each  25  cm8  0.0801  grm.  of 
HC1,  agreeing  with  0.0794  grm.  and  0.0791  grm.  obtained  by 
the  new  method.  An  analysis  of  the  sulphuric  acid  solution 


ANALYSIS  OF  ALKALIES  AND  ACIDS.  171 

by  the  Groger  method  gave  for  each  25  cm3  0.1241  grm.  of 
H2SO4  agreeing  with  0.1246  grm.  obtained  by  the  new  method. 
This  investigation  shows  that  the  reaction  between  iodine 
and  hydroxides  of  the  alkalies  and  alkaline  earths  in  hot 
solution  is  regular  and  complete  under  analytical  conditions, 
not  being  appreciably  affected  by  the  mass  action  of  considerable 
excesses  of  iodine.  The  reaction  is  best  applied  in  analysis  by 
titrating  the  alkali  with  an  excess  of  iodine,  removing  this 
excess  by  boiling,  and  estimating  the  iodine  in  the  residue. 
While  certain  mechanical  difficulties  may  effect  the  extreme 
accuracy  of  the  process  as  a  direct  means  for  analyzing  alkalies, 
the  action  is  at  all  times  regular  and  may  be  indirectly  applied 
with  fair  accuracy  to  the  analysis  of  various  acids  and  possibly 
to  other  compounds.  The  reaction  between  iodine  and  alkali 
carbonates  on  the  contrary  is  irregular  and  cannot  be  made 
the  basis  of  any  analytical  process. 


XXII 

THE  ESTIMATION  OF  BORIC  ACID. 

BY  F.  A.  GOOCH  AND  LOUIS  CLEVELAND  JOKES.* 

THE  estimation  of  boric  acid  by  treating  the  salts  of  that 
acid  with  sulphuric  acid,  distilling  with  methyl  alcohol, 
evaporating  the  distillate  over  magnesium  oxide,  igniting  and 
weighing,  was  proposed  by  Rosenbladt-t  A  little  later,  and 
without  knowledge  of  Rosenbladt's  experience,  a  somewhat 
similar  process,  J  which  consisted  in  the  treating  of  the 
compound  of  boric  acid  with  acetic  acid  or  nitric  acid, 
distillation  with  methyl  alcohol,  evaporation  of  the  distillate 
over  calcium  oxide,  and  ignition  of  the  residue,  was  described 
by  one  of  us.  In  the  course  of  the  development  of  this  process, 
it  transpired  that  the  insolubility  of  magnesium  oxide  retards 
the  absorption  of  boric  acid  by  that  substance,  and  that  the 
more  soluble  calcium  oxide  retains  boric  acid  more  actively  and 
is  therefore  to  be  preferred. 

Points  in  the  treatment  upon  which  special  emphasis  was 
laid  in  the  original  description  of  this  process  were  the  choice 
of  a  suitable  apparatus  for  the  distillation,  the  employment  of 
a  loosely  stoppered  receiver  for  the  reception  of  the  distillate 
upon  slaked  lime,  the  careful  removal  of  water  from  the 
substance  in  the  retort  before  acidifying  and  treating  with  the 
methyl  alcohol,  regulated  use  of  acid,  and  care  in  the  evapora- 
tion and  ignition. 

The  attainment  of  good  results  in  this  process  depends  upon 
attention  to  details.  Modifications  have  been  suggested  by 
several  investigators.  Thus,  instead  of  igniting  the  calcium 
oxide  in  a  large  platinum  crucible,  transferring  it  to  the 

*  From  Am.  Jour.  Sci.,  vii,  34.  t  Zeitschr.  anal.  Chem.,  xxvi,  21. 

J  Am.  Chem.  Jour.,  ix,  23. 


ESTIMATION  OF  BORIC  ACID.  173 

receiver  to  hold  the  boric  acid,  and  returning  the  calcium  oxide 
with  the  distillate  to  the  same  crucible  for  subsequent  ignition 
of  the  residue,  as  was  originally  proposed,  Penfield*  prefers 
to  ignite  the  calcium  oxide  in  a  small  crucible,  to  collect  the 
distillate  hi  ammoniacal  water,  to  evaporate  the  latter  over  the 
calcium  oxide  in  a  large  platinum  dish,  and  to  transfer  this 
residue  back  to  the  small  crucible  for  the  final  evaporation  and 
ignition.  Kraut  f  suggests  a  modification  of  form  in  the 
apparatus  with  no  other  essential  change  hi  conditions. 
MoissanJ  has  suggested  changes  in  the  apparatus  and  avoids 
a  transfer  of  the  calcium  oxide — collecting  the  distillate  by 
itself  in  a  closed  receiver,  trapped  with  an  ammonia  bulb  to 
prevent  the  escape  of  the  boric  acid  from  the  distillate; 
furthermore,  Moissan's  process  calls  for  the  use  of  an  amount 
of  calcium  oxide  from  fifteen  to  twenty  tunes  greater  than 
that  theoretically  required.  From  our  experience  it  seems 
obvious  that  the  demand  for  this  amount  of  calcium  oxide 
arises  from  an  excessive  use  of  nitric  acid  in  the  retort  and  the 
consequent  modification  of  conditions  in  the  distillate.  For- 
tunately this  difficulty  may  be  avoided  by  the  use  of  a  little 
phenolphthalein  as  an  indicator  in  the  retort  and  care  to  limit 
the  addition  of  nitric  acid  to  the  amount  required  to  produce 
distinct  acidity.  The  addition  of  a  drop  of  the  acid  and 
another  of  the  indicator  should  be  repeated  once  or  twice 
during  the  distillation  to  insure  the  return  of  the  volatilized 
acid  to  the  salt  slightly  decomposed  in  the  process.  The 
effect  of  much  nitric  acid  is  bad,  not  only  because  it  neu- 
tralizes the  calcium  oxide  when  it  passes  to  the  distillate, 
but  because  when  it  is  used  a  tendency  is  developed  on  the 
part  of  the  dried  mixture  of  calcium  hydroxide  and  borate  to 
puff  explosively  if  the  ignition  is  begun  as  soon  as  the  residue 
is  dry.  If  the  residue  is  heated  gradually  and  as  strongly  as 
possible  over  a  radiator  before  the  flame  is  actually  applied 
to  the  crucible,  no  such  action  takes  place ;  we  are  disposed  to 
attribute  it  to  the  effect  of  the  nitrate  and  nitrite,  produced  by 

*  Am.  Jour.  Sci.,  xxxiv,  222.  t  Zeitschr.  anal.  Chem.,  xxxvi,  3. 

t  Comp.  rend.,  cxvi,  1084. 


174 


ESTIMATION  OF  BORIC  ACID. 


the  absorption  of  nitrous  fumes  in  the  lime,  upon  the  alcohol 
or  other  organic  matter  retained  by  the  lime  hi  the  evaporation 
and  drying  unless  the  latter  process  is  prolonged  at  high 
temperature. 

That  good  results  may  be  obtained  with  small  amounts  of 
calcium  oxide,  provided  care  as  to  the  use  of  nitric  acid  and 
the  conditions  of  ignition  be  taken,  is  shown  by  the  figures 
of  the  original  description  and  by  the  following  experiments, 
in  which  phenolphthalein  was  employed  as  an  indicator  and 
the  residue  heated  strongly  over  the  radiator  before  actual 
ignition. 


CaO  taken. 

B2O8  taken. 

B203  found. 

Error. 

grin. 

2.3405 
1.7620 
2.1757 
2.5656 

grm. 
0.1788 
0.1790 
0.1824 
0.1788 

grm. 
0.1792 
0.1785 
0.1840 
0.1786 

grm. 
0.0004+ 
0.0005- 
0.0016+ 
0.0002- 

These  results  are  accurate  within  reasonable  limits.  On 
the  other  hand,  without  care  to  ignite  gradually  we  have 
noted  errors  of  from  0.0030  grm.  to  0.0060  grm.  in  the  process 
otherwise  conducted  similarly.  Doubtless  the  use  of  large 
amounts  of  calcium  oxide  as  suggested  by  Moissan  may  serve 
the  purpose  of  diffusing  the  explosive  mixture  through  a 
mass  of  inert  matter  sufficient  to  prevent  violent  puffing,  but 
care  to  heat  over  the  radiator  as  strongly  as  possible  before 
opening  the  flame  directly  to  the  crucible  answers  the  same 
end.  The  difficulty  does  not  exist  when  acetic  acid  is  used 
in  place  of  nitric  acid,  though  even  in  this  case  it  is  safer  to 
use  the  radiator  in  the  first  stages  of  heating,  thus  avoiding 
the  danger  of  mechanical  loss  by  too  rapid  ignition. 

Following  are  determinations  made  by  this  method  with  the 
use  of  acetic  acid.  The  results  of  these  experiments,  as 
well  as  those  of  the  investigators  mentioned,  are  a  sufficient 
answer  to  the  criticism  of  Reischle,*  that  acetic  acid  and 
nitric  acid  do  not  liberate  boric  acid  in  the  distillation  pro- 
*  Zeitschr.  anal.  Chem.,  xxvi,  512. 


ESTIMATION  OF  BORIC  ACID. 


175 


CaO  taken. 

B203  taken. 

B,O3  found. 

Error. 

gnu. 
0.9977 
1.0220 
1.3717 
1.1310 

gnu. 
0.2065 
0.2067 
0.2077 
0.1791 

gnu. 
0.2062 
0.2070 
0.2075 
0.1795 

gnu. 
0.0003- 
0.0003+ 
0.0002- 
0.0004-1- 

cess  so  that  good  results  may  be  obtained.  Moreover,  it  has 
been  shown  by  one  of  us  *  that  even  carbonic  acid  is  strong 
enough  to  bring  about  complete  volatility  of  boric  acid  with 
methyl  alcohol. 

The  use  of  Calcium  Oxide  as  a  Retainer. 

Quite  recently  Thaddeeff  f  has  advocated  the  abandonment 
of  calcium  oxide  as  an  agent  for  holding  boric  acid  in  the 
evaporation  of  alcoholic  and  aqueous  solutions,  on  account 
of  the  hygroscopic  nature  of  the  oxide  and  the  consequent 
difficulty  of  securing  it  in  definite  conditions  for  weighing, 
and  proposes,  instead  of  using  calcium  oxide,  to  retain  and 
estimate  boric  acid  in  solution  by  converting  it  into  the  form 
of  potassium  borofluoride. 

In  the  final  modification  of  ThaddeefFs  method  the  proposal 
is  made  to  liberate  the  boric  acid  from  its  compounds  by 
sulphuric  acid,  to  volatilize  it  in  methyl  alcohol  with  the 
aid  of  a  current  of  dry  air,  to  catch  the  distillate  in  potassium 
hydroxide,  to  treat  the  mixture  of  hydroxide  and  borate  with 
hydrofluoric  acid  in  excess  and  evaporate  on  the  steam  bath, 
to  digest  the  residue  of  fluoride  and  borofluoride  at  normal 
temperatures  for  two  hours  with  50  cm3  of  a  potassium 
acetate  solution  (sp.  gr.  1.14)  and  for  twelve  hours  more  after 
adding  100  cm3  of  ethyl  alcohol  (sp.  gr.  0.805),  to  filter  on 
paper,  wash  the  residue  with  62-72  cm3  of  alcohol  (sp.  gr. 
0.805),  dry  at  100°  and  weigh  as  potassium  borofluoride, 
after  which  the  borofluoride  is  to  be  dissolved  in  boiling  water 
and  tested  with  calcium  chloride  for  possible  contamination 

*  Jones,  Am.  Jour.  Sci.,  v,  442.    This  volume,  p.  100. 
t  Zeitschr.  anal.  Chem.,  xxxvi,  568. 


176  ESTIMATION  OF  BORIC  ACID. 

by  the  presence  of  a  fluoride.  Plainly  ThaddeefFs  procedure 
presents  at  the  outset  difficulties;  for  besides  the  incon- 
venience of  conducting  long  digestions  with  reagents  of 
regulated  strength,  the  difficulty  of  procuring  hydrofluoric 
acid  free  from  silica,  which  if  present  (as  it  usually  is  hi  the 
so-called  chemically  pure  hydrofluoric  acid  of  commerce) 
would  be  retained  in  the  borofluoride  as  potassium  fluosilicate, 
the  inaccuracy  of  the  dried  paper  filter,  and  the  obvious 
uncertainty  of  success  in  an  attempt  to  wash  a  mixture  of 
acid  potassium  fluoride  and  potassium  borofluoride  in  potas- 
sium acetate  and  alcohol  so  that  the  one  shall  be  rendered 
entirely  soluble  while  the  other  remains  sensibly  unaffected, 
—  besides  these  objections,  there  is  the  theoretical  probability 
that  boric  acid  mnst  be  lost  by  volatilization  during  the 
evaporation  of  the  solution  of  the  mixed  salts  in  the  presence 
of  free  hydrofluoric  acid.  This  last  point  was  put  to  the 
proof  by  submitting  to  distillation  in  a  platinum  retort  a 
mixture  of  equal  quantities  of  borax  and  potassium  hydroxide 
with  an  excess  of  hydrofluoric  acid,  collecting  the  distillate 
in  potassium  hydroxide,  evaporating  it  to  dryness  and  testing 
it  for  the  presence  of  boric  acid.  When  this  residue  from  the 
evaporated  distillate  was  treated  with  sulphuric  acid  and 
methyl  alcohol,  the  burning  alcohol  vapor  gave  plainly 
the  green  flame  of  boric  acid.  Another  portion  showed 
clearly  the  presence  of  boric  acid  when  acidulated  with 
hydrochloric  acid  tested  with  turmeric  paper.  No  boric  acid 
could  be  detected  in  any  of  the  reagents  used.  It  is  plain, 
therefore,  that  boric  acid  does  volatilize  upon  the  evapora- 
tion of  a  mixture  of  potassium  fluoride  and  borofluoride  in 
acid  solution.  The  amount  of  such  loss  is  disclosed  hi  the 
record  of  the  following  experiment.  Portions  of  a  standard 
solution  of  boric  acid,  prepared  by  dissolving  a  known  weight 
of  anhydrous  boric  oxide  in  a  liter  of  water,  were  mixed 
with  a  solution  of  potassium  hydroxide  (free  from  silica  and 
standardized  by  conversion  to  the  chloride)  in  the  proportions 
to  form  the  potassium  borofluoride,  and  an  excess  of  hydro- 
fluoric acid  was  added.  The  mixture  was  evaporated  and 


ESTIMATION  OF  BORIC  ACID. 


177 


the  residue  was  dried  and  weighed  at  100°,  the  whole  opera- 
tion being  conducted  in  platinum. 


Exp. 

HKP, 
equivalent  to 
KOH  taken. 

taken. 

KPBFs 
theoretical 
weight. 

KFBF8 
found. 

Error  in 
terms  of 
KFBF8. 

Error  in 
terms  of 
8,03. 

grm. 

grin. 

grm. 

grin. 

grm. 

grm. 

(1) 

0.3531 

0.1582 

0.5701 

0.5580 

0.0121- 

0.0033- 

(2) 

0.3192 

0.1430 

0.5154 

0.5100 

0.0054- 

0.0015- 

(3) 

0.3192 

0.1430 

0.5154 

0.5030 

0.0124- 

0.0034- 

(4) 

0.3192 

0.1430 

0.5154 

0.5088 

0.0066- 

0.0018- 

(6) 

0.3192 

0.1430 

0.5154 

0.5114 

0.0040- 

0.0011- 

In  experiments  (1)  to  (3)  the  volume  of  the  solution 
evaporated  was  about  50  cm3.  In  experiment  (4)  this 
volume  was  reduced  about  one-half  before  acidifying  with 
hydrofluoric  acid,  while  in  experiment  (5)  the  solution  was 
diluted  about  one-half  before  adding  the  hydrofluoric  acid. 
It  is  plain,  therefore,  that  in  this  single  step  of  Thaddeeff's 
process  there  is  a  considerable  error  of  deficiency.  On  the 
other  hand,  the  errors  for  the  full  process  as  laid  down  by 
Thaddeeff  have  been  in  our  experience  invariably  differences 
of  excess  —  presumably  because  the  loss  due  to  volatilization 
of  boric  acid  has  been  overbalanced  by  the  inaccuracy  in 
washing.  It  is  plain  that  the  process  can  give  true  indica- 
tions only  by  the  balancing  of  considerable  errors. 

If  we  take  into  consideration,  therefore,  the  inevitable 
inaccuracy  and  inconvenience  of  Thaddeeff's  proposal,  it 
cannot  be  regarded  as  a  desirable  substitute  for  the  process 
according  to  which  boric  acid  is  absorbed  and  retained  for 
weighing  with  calcium  oxide,  especially  since  the  difficulties 
in  the  way  of  getting  constant  weights  of  that  substance  are 
by  no  means  insuperable. 

Thus  the  following  table  shows  the  series  of  weights  taken 
in  several  experiments  in  bringing  calcium  oxide  to  a  constant 
weight  in  a  50  cm3  platinum  crucible  ignited  over  a  blast  lamp, 
as  well  as  the  weight  taken  after  adding  a  known  amount  of 
standard  boric  acid  solution  to  the  slaked  oxide,  evaporating, 
and  igniting.  The  results  recorded  are  those  of  experiments 

VOL.   II. — 12 


178 


ESTIMATION  OF  BORIC  ACID. 


made  on  days  not  moist  beyond  the  average  and  with  the 
greatest  care  to  approach  the  limit  of  accuracy  with  which 
calcium  oxide  and  the  boric  acid  held  thereby  can  be  weighed 
under  ordinarily  favorable  conditions.  The  first  weight  of 
calcium  oxide  recorded  under  each  experiment  was  taken  after 
a  strong  ignition  over  the  blast  lamp  for  about  one-half  hour. 
The  succeeding  weights  were  taken  after  similar  ignitions  of 
five  minutes.  In  all  cases  the  crucible  was  left  to  stand  a 
definite  period  in  a  sulphuric  acid  desiccator,  and,  after  the 
approximate  value  had  once  been  obtained,  the  weights  of 
the  preceding  weighing  were  replaced  on  the  balance  before 
the  crucible  was  taken  from  the  desiccator.  The  average  of  the 
weights  bracketed  is  the  weight  taken  as  constant  for  the 
calculations. 


Exp. 

CaO  taken. 

B,08 
taken. 

CaO  +  B208 
taken. 

CaO  +  B2O8 
found. 

Error. 

(1) 

grm. 
0.9505 
0.9493  I  ft  0400 
0.9493  f  U<W 

grm. 

0.2095 

grm. 

1.1588 

grin. 

1.1590  )  i  1591 

1.1591  f  L1591 

grm. 

0.0003+ 

(2) 

1.1319 
1.1317  ) 
1.1313  }  1.1315 
1.1315  ) 

0.2150 

1.3465 

1.3499 
1.3474  ) 
1.3475  }  1.3475 
1.3476  ) 

0.0010+ 

(3) 

0.8028 
0.8025)  ft  o0o4 
0.8024  f0'8024 

0.1184 

0.9208 

0.9205 
0.9206  )  ft  oonft 
0.9206  f0'9206 

0.0002- 

(4) 

2.6980 
2.6975 
2.6973  )  9  fiQ7o 
2.6973  I2 

0.2073 

2.9046 

2.9043 
2.9049  Uqo4R 
2.9044  ]  ^>yu4b 

0.0002+ 

Obviously  calcium  oxide  may  be  weighed  with  accuracy, 
with  or  without  boric  acid ;  but  the  fact  remains  that  a  less 
hygroscopic  absorbent  —  one  requiring  less  care  in  the  handling, 
is  to  be  desired. 

The  use  of  Sodium  Tungstate  as  a  Retainer. 

In  searching  for  a  suitable  material  of  less  hygroscopicity  to 
replace  calcium  oxide  as  a  retainer  for  boric  acid,  we  have 
found  that  sodium  tungstate,  fused  with  a  slight  excess  of 


ESTIMATION  OF  BORIC  ACID. 


179 


tungstic  acid  over  that  contained  in  the  normal  tungstate 
(to  insure  its  freedom  from  carbonate),  answers  this  purpose 
excellently.  This  substance  is  definite  in  weight,  not  hygro- 
scopic, soluble  in  water,  and  recoverable  in  its  original  weight 
after  evaporation  and  ignition.  To  test  its  value  as  a  retainer 
for  boric  acid,  portions  of  it  —  4  to  7  grm.  —  were  fused  and 
weighed  in  a  50  cm3  crucible,  the  tungstate  was  dissolved  in 
water  and  to  it  was  added  a  known  amount  of  a  standard 
solution  of  boric  acid.  After  diluting,  mixing,  evaporating, 
and  fusing  the  residue,  the  increase  in  weight  should  represent 
the  boric  anhydride  held  by  the  tungstate.  The  results  of  the 
accompanying  table  show  how  accurately  the  boric  acid  is 
retained  under  these  conditions.  In  experiments  (3)  to  (7) 
the  tungstate,  after  its  first  weighing,  was  dissolved,  transferred 
to  a  larger  platinum  dish  and  mixed  therein  with  the  boric 
acid.  After  evaporation  to  a  suitable  volume  this  solution  of 
tungstate  and  boric  acid  was  transferred  to  the  original  crucible 
for  final  evaporation  and  ignition. 


Exp. 

NajWOi  +  W03 
taken. 

B208  taken. 

B,03  found. 

Error  in  B203. 

grm. 

grm. 

grm. 

grm. 

(1) 

6.6416 

0.1784 

0.1771 

0.0013- 

(2) 

7.3134 

0.1786 

0.1773 

0.0013- 

(3) 

5.5003 

0.0950 

0.0952 

0.0002-f- 

(4) 

4.1394 

0.0944 

0.0944 

0.0000 

(5) 

7.5037 

0.2148 

0.2149 

0.0001+ 

(6) 

4.7744 

0.2718 

0.2702 

0.0016- 

(7) 

6.6470 

0.2503 

0.2487 

0.0016- 

It  is  plain  that  though  the  sodium  tungstate  does  not  hold 
the  boric  acid  with  absolute  accuracy  the  errors  are  not 
unreasonable  —  0.0008  grm.  in  the  mean.  Upon  substituting 
the  tungstate  for  calcium  oxide  as  a  retainer  in  the  distillation 
process,  the  results  were  likewise  highly  favorable. 

We  used  by  preference  the  apparatus  originally  proposed, 
excepting  that  the  Erlenmeyer  flask  used  as  a  receiver  is  fitted 
tightly  to  the  condenser  and  trapped  with  water  bulbs.  The 
retort  is  made  very  easily  from  a  150  cm3  pipette  and  has 


180 


ESTIMATION  OF  BORIC  ACID. 


the  special  advantage  that  particles  of  the  residue  spattering 
during  distillation  are  easily  washed  from  the  walls  of  the 
vessel  by  a  slight  rotary  motion  of  the  retort.  It  was  found 
that  special  care  should  be  taken  to  give  the  tungstate  ample 
time  for  contact  with  the  distillate  before  exposing  the  latter 
to  atmospheric  evaporation.  The  distilkte  was  received, 
therefore,  in  a  dilute  solution  of  sodium  tungstate  placed  in 
the  receiver,  cooled  by  ice  and  trapped  with  water,  and  the 
mixture  was  well  stirred,  allowed  to  stand  one  half-hour, 
evaporated  to  small  volume  in  a  large  dish,  and  transferred 
to  the  crucible  in  which  the  tungstate  had  been  originally 
weighed.  After  thorough  drying  the  residue  was  ignited  to 
fusion  and  weighed.  When  acetic  acid  was  employed  in  the 
retort,  care  was  taken  in  the  ignition  to  expose  the  fused  mass 
freely  to  the  air  (by  causing  it  to  flow  upon  the  sides  of  the 
crucible)  until  the  color  of  the  cooled  tungstate  was  white,  in 
order  that  the  reducing  effect  of  the  acetate  might  be  eliminated. 
In  the  experiments  recorded  in  the  following  table  the  tungstate 


Na,W04  +  WOS 

B,03  taken. 

B203  found. 

Error. 

WITH  NITRIC  ACID. 

gnu. 

8.5516 
4.9639 
8.0033 

grin. 

0.1582 
0.1329 
0.1267 

grm. 

0.1572 
0.1323 
0.1256 

grm. 

0.0010- 
0.0006- 
0.0011- 

WITH  ACETIC  ACID. 

4.9658 
6.0289 
4.6797 
4.0013 

0.1434 
0.1431 
0.1589 
0.1433 

0.1418 
0.1433 
0.1587 
0.1422 

0.0016- 
0.0002+ 
0.0002- 
0.0011- 

WITH  SULPHURIC  ACID. 

6.3439 
8.8227 
10.1516 
6.5738 

0.1582 
0.1582 
0.1265 
0.1392 

0.1579 
0.1577 
0.1264 
0.1390 

0.0003- 
0.0005- 
0.0001- 
0.0002- 

was  used  in  the  receiver  to  retain  the  boric  acid  distilled  as 
usual  with  methyl  alcohol,  from  the  borates  treated  with  acetic 


ESTIMATION  OF  BORIC  ACID.  181 

acid,  nitric  acid  or  sulphuric  acid,  in  amounts  regulated  by  the 
use  of  phenolphthalein  as  an  indicator. 

Excessive  use  of  acid  is  disadvantageous,  and  this  is 
especially  true  in  the  case  of  sulphuric  acid;  for,  if  this 
acid  is  carried  over  with  the  methyl  alcohol,  as  it  is  at  100° 
if  present  in  appreciable  excess,  a  part  of  it,  at  least,  is  held 
permanently  by  the  tungstate  to  increase  the  apparent  weight 
of  the  boric  acid  to  be  estimated. 

The  manipulation  of  the  tungstate  presents  no  difficulties, 
and  the  results  obtained  by  its  use  are  reasonably  accurate. 


XXIII 

A  VOLUMETRIC  METHOD  FOR  THE  ESTIMATION 
OF  BORIC  ACID. 

BY  LOUIS  CLEVELAND  JONES.* 

WHEN  boric  acid  and  mannite  are  mixed  in  solution  a  peculiar 
compound  of  strongly  acid  properties  is  the  result.  This  com- 
pound decomposes  carbonates,  and  its  acid  taste  is  comparable 
to  that  of  citric  acid,  much  stronger  than  that  of  boric  acid 
alone.  Magnaninit  has  found  that  the  product  of  such  a 
mixture  of  boric  acid  and  mannite  solutions  shows  greater 
electrical  conductivity  and  a  lower  freezing  point  than  a  simi- 
lar molecular  solution  of  either  substance  alone.  Other  poly- 
atomic alcohols  (but  all  to  a  less  degree  than  mannite)  and 
some  organic  acids  show  this  peculiar  property  of  combining 
chemically  with  boric  acid  to  increase  its  acid  qualities.^:  Of 
this  reaction  between  boric  acid  and  the  polyatomic  alcohols, 
Thomson,§  Barthe,||  and  Jb'rgensen^"  have  taken  advantage 
to  develop  methods  for  the  volumetric  estimation  of  boric  acid. 
Glycerine  is  used  to  form  a  combination  with  boric  acid,  suffi- 
ciently acidic  to  give  an  acid  reaction  when  used  with  a  sensi- 
tive indicator  and  make  possible  its  titration  with  an  alkaline 
solution.  Honig  and  Spitz**  show  that  in  the  method  of 
Jb'rgensen  a  very  large  amount  of  glycerine  must  be  used  to 
prevent  the  appearance  of  the  indication  of  alkalinity  with  phe- 
nolphthalein  before  all  the  boric  acid  is  neutralized  according 
to  the  following  equation,  2NaOH+B2O3  =  2NaOBO+H2O ; 

*  From  Am.  Jour.  Sci.,  vii,  147. 
t  Gaz.  Chim.,  xx,  428-440;  xxi,  134-145. 

t  Klein,  Jour.  Pharm.  Chim.,  4,  vol.  xxviii ;  Lambert,  Comp.  rend.,  cviii, 
1016-1017. 

§  Jour.  Soc.  Chem.  Ind.,  xv,  432.          ||  Jour.  Pharm.  Chim.,  xxix,  163. 
IT  Zeitschr.  angew.  Chem.,  1897, 5.     **  Zeitschr.  angew.  Chem.  (1896),  649. 


ESTIMATION  OF  BORIC  ACID.  183 

that  in  the  presence  of  carbonates  the  solution  must  be  boiled 
to  decompose  bicarbonates  and  the  escape  of  boric  acid  by  vol- 
atilization prevented  by  the  use  of  a  return  condenser;  and 
that  silica  must  be  removed  by  the  process  of  Berzelius,  and 
the  solution  then  neutralized  by  the  use  of  methylorange 
before  a  titration  of  the  boric  acid  can  be  made. 

Vadam,*  for  the  estimation  of  boric  acid  in  butter  makes 
use  of  mannite,  which,  as  he  finds,  gives  sharper  indication 
with  litmus  than  glycerine.  According  to  this  process,  the 
solution  to  be  analyzed  for  boric  acid  is  neutralized  by  the  use 
of  litmus  and  a  solution  of  sodium  hydroxide.  Mannite 
(1-2  grm.)  is  then  added,  bringing  about  an  acid  reaction  with 
the  boric  acid  present  in  free  condition.  The  solution  is  then 
titrated  to  alkalinity  by  sodium  hydroxide. 

None  of  the  above  methods  with  glycerine  have,  in  my 
experience,  given  anything  but  comparatively  crude  results. 
The  weak  acidic  properties  of  boric  acid,  the  interference  (and 
difficulty  of  removal)  of  carbon  dioxide  with  all  organic  indi- 
cators sufficiently  delicate  to  be  used  with  boric  acid,  and 
indeed,  the  procuring  of  a  standard  alkali  containing  no  car- 
bonate, together  with  the  supposed  detrimental  influence  of 
silica  and  the  lack  of  a  convenient  method  for  its  removal, 
have  made  the  process  of  Gooch,f  which  involves  distillation 
and  weighing  with  calcium  oxide,  the  only  means  (though 
requiring  long  time  and  exceeding  care)  in  use  for  the  accu- 
rate separation  and  estimation  of  boric  acid.  Recently  sodium 
tungstate  has  been  recommended  from  this  laboratory^  as 
a  substitute  for  calcium  oxide  to  retain  the  distilled  boric  acid. 
The  entire  process,  however,  is  one  of  the  most  exacting  in 
analytical  chemistry,  and  for  this  reason  a  convenient,  rapid 
and  at  the  same  time  accurate  method  for  the  estimation  of 
boron  is  especially  desirable.  The  first  step  toward  the 
development  of  such  a  process  must  be  the  convenient  prepa- 

*  Jour.  Pharm.  Chim.  (6),  viii,  109-111. 

t  Am.  Chem.  Jour,  ix,  23-33;  Moissan,  Comp.  rend.,  cxri,  1087;  Kraut, 
Zeitschr.  anal.  Chem.  xxxvi,  165 ;  Montemartini,  Gaz.  Chim.  Ital.,  xxviii,  1, 
344. 

t  Gooch  and  Jones,  Am.  Jour.  Sci.,  rii,  34.    This  volume,  p.  172. 


184  A    VOLUMETRIC  METHOD  FOR    THE 

ration  and  the  accurate  estimation  of  the  standard  solution  of 
alkali  to  be  used  for  neutralizing  the  boric  acid.  This  has  been 
found  to  be  easily  accomplished  by  the  process  recommended 
by  Kusler.*  This  observer,  in  an  extensive  investigation  of 
the  analytical  methods  for  the  volumetric  estimation  of  alkalies 
and  alkali  carbonates  in  solution,  finds  that  both  phenolphthal- 
ein  and  methylorange  are  appreciably  sensitive  to  carbonic  acid, 
but  when  this  interfering  agent  is  removed  by  precipitation 
with  barium  chloride  according  to  the  process  of  Winkler,f  the 
remaining  free  alkali  may  be  estimated  with  great  accuracy  by 
phenolphthalein  and  decinormal  hydrochloric  acid. 

Obviously,  if  the  difficulties  dependent  upon  the  action  of 
carbon  dioxide  can  be  obviated,  and  if  the  acidity  of  the  boric 
acid  can  be  increased  to  such  an  extent  that  a  sufficiently  sen- 
sitive indicator  will  give  with  accuracy  the  neutralization 
point  with  free  alkali,  and  if  the  alkali  and  stronger  acid  can 
be  combined  while  boric  acid  alone  remains  free,  then  it 
should  be  possible  to  estimate  boric  acid  volumetrically.  Ex- 
periment has  shown  that  barium  chloride  removes  carbon 
dioxide  present  in  carbonates,  and  that  mannite  makes  a  com- 
bination with  boric  acid  strongly  acidic  to  phenolphthalein. 

To  obtain  the  boric  acid  alone  in  free  condition  many 
attempts  have  been  made.  Gladding,f  Thaddeeff  §  and  Rosen- 
bladt  ||  have  isolated  the  boric  acid  by  distillation  with  methyl- 
alcohol  and  a  non-volatile  acid.  Many  indicators  theoretically 
insensible  to  free  boric  acid  have  been  used  to  indicate  the 
neutralization  of  the  stronger  acids.  Honig  and  Spitz,*[f  and 
Thomson,**  use  methylorange,  Morse  and  Burton,|t  tropae- 
olin  00,  while  Vadam  JJ  makes  use  of  litmus.  All  these 
indicators,  however,  have  been  found  by  experiment  to  be 
more  or  less  affected  by  boric  acid  in  solution.  On  the  other 
hand,  I  have  found  in  the  well  known  reaction  according 

*  Zeitschr.  anorg.  Chem.,  xiii,  124-150.  t  Massanalyse. 

t  Jour.  Am.  Chem.  Soc.,  iv,  568. 

§  Zeitschr.  anal.  Chem.,  xxxvi  (9),  568. 

||  Zeitschr.  anal.  Chem.,  xxvi,  18.         f  Zeitschr.  anorg.  Chem.  (18),  549. 
**  Jour.  Soc.  Chem.  Ind.,  xv,  432.          ft  Am.  Chem.  Jour.,  x,  154. 
|J  Jour.  Pharm.  Chim.  (6)  viii,  109-111. 


ESTIMATION  OF  BORIC  ACID.  185 

to  which  a  stronger  acid  liberates  regularly  iodine  from  a 
mixture  of  iodide  and  iodate,  the  solution  of  this  difficulty. 
If  both  the  iodide  and  iodate  are  in  excess  of  the  acid  the 
entire  amount  of  free  acid  will  be  neutralized  and  the  cor- 
responding amount  of  iodine  liberated  according  to  the 
following  equation: 

5KI  +  KI03  +  6HC1  =  6KC1  +  3H20  +  3I2. 

This  liberated  iodine  may  be  removed  by  sodium  thiosulphate 
and  a  solution  obtained  which  is  absolutely  neutral  containing 
only  neutral  salts,  potassium  iodide,  iodate,  and  tetrathionate. 
The  statements  made  by  P.  George  vie*  and  Furry,  f  that 
boric  acid  present  in  moderate  amount  in  solution  has  not 
the  slightest  action  on  a  mixture  of  iodide  and  iodate,  have 
been  experimentally  verified.  Therefore  when  this  acid  is 
liberated  by  an  excess  of  a  stronger  acid,  and  the  iodine  set 
free  destroyed  by  thiosulphate,  it  remains  free  in  solution  to 
be  titrated  in  any  convenient  manner  possible. 

Following  along  the  lines  suggested  by  the  above  reac- 
tions, a  volumetric  process  for  the  estimation  of  boric  acid 
has  been  developed.  For  a  basis  of  the  investigations,  a 
standard  solution  of  boric  acid  was  prepared  by  dissolving 
in  a  liter  of  water  about  eight  grams  of  carefully  weighed 
anhydrous  boric  oxide.  This  anhydrous  boric  oxide  was 
prepared  from  the  several  times  recrystallized  hydrous  boric 
acid,  by  long-continued  fusion  over  a  blast  lamp.  A  solution 
of  approximately  \  sodium  hydroxide  was  prepared  from  the 
ordinary  sodium  hydroxide  of  the  laboratory.  The  free  alkali 
in  this  solution  was  estimated  by  the  process  of  Winkler 
recommended  by  Ktisler.  The  acid  used  to  make  this  estima- 
tion was  hydrochloric,  standardized  by  silver  nitrate. 

The  full  method  for  the  estimation  of  boric  acid  as  finally 
elaborated  is  as  follows :  The  solution  is  made  clearly  acid 
to  litmus  by  hydrochloric  acid  and  5  cm3  of  a  solution  (10%) 
of  barium  chloride  added.  An  amount  of  iodate  and  iodide 
of  potassium  sufficient  to  liberate  an  amount  of  iodine  at 

*  Jour,  prakt.  Chem.,  xxxviii,  118.  t  Am.  Chem.  Jour.,  vi,  341. 


186  A    VOLUMETRIC  METHOD  FOR   THE 

least  equivalent  to  the  excess  of  hydrochloric  acid  in  the 
acidified  solution  is  mixed  with  starch  in  a  separate  beaker, 
and  the  iodine  which  is  usually  thrown  out  by  this  mixture, 
is  just  bleached  by  a  dilute  solution  of  thiosulphate. 

To  the  now  neutral  solution  of  iodide  and  iodate  a  single 
drop  of  the  solution  to  be  analyzed  is  transferred  by  a  glass 
rod.  If  a  blue  coloration  is  developed,  the  solution  is  acidic 
with  hydrochloric  acid,  and  all  the  boric  acid  is  in  free 
condition.  The  amount  of  iodide  and  iodate  used  depends 
upon  the  acidity  of  the  solution  containing  boric  acid. 
Usually  10  cm3  of  a  25  per  cent  solution  of  iodide  and  the 
same  amount  of  a  saturated  solution  of  iodate  is  sufficient. 
Any  larger  excess  of  hydrochloric  acid  should  be  neutralized 
by  sodium  hydroxide  before  the  iodide  and  iodate  mixture  is 
added.  After  the  addition  of  the  iodide  and  iodate  solution, 
containing  starch,  to  the  boric  acid  solution,  the  liberated 
iodine  should  be  carefully  bleached  by  thiosulphate.  Any 
excess  of  thiosulphate  in  reasonable  amount  does  not  seem 
to  be  detrimental,  but  in  practice  the  starch  iodide  color  is 
clearly  bleached,  and  no  more  then  added.  Carbonates  pre- 
vent a  definite  indication  of  the  neutral  point  by  thiosulphate 
and  starch  iodide,  therefore  the  barium  chloride  (about  5  cm3) 
should  be  added  before  this  point  in  the  process.  The 
mixture  of  iodide  and  iodate  is  not  added  to  the  solution  to 
be  analyzed  until  after  it  is  made  acidic,  for  the  reason  that 
when  the  neutral  point  is  approached  by  the  addition  of 
hydrochloric  acid  the  starch  iodide  is  thrown  out  locally  by 
the  acid,  and  the  small  amount  of  sodium  borate  remaining 
undecomposed  does  not  again  bleach  the  coloration  produced 
thus  obscuring  the  neutral  point  which  must  be  obtained 
before  titrating  for  boric  acid. 

The  solution  after  the  bleaching  of  the  iodine  by  thiosul- 
phate is  colorless  and  contains  only  starch,  neutral  chloride, 
potassium  tetrathionate,  iodide  and  iodate,  and  all  the  boric 
acid  present  in  uncombined  condition.  The  carbonate  lies  out 
of  the  sphere  of  action  in  insoluble  form  as  barium  carbonate. 
A  few  drops  of  the  indicator,  phenolphthalein,  are  now  added, 


ESTIMATION  OF  BORIC  ACID.  187 

and  the  alkaline  solution  run  in  until  a  strong  red  coloration  is 
produced.  A  pinch  of  mannite  is  then  added,  which  bleaches 
the  phenolphthalein  coloration,  and  the  alkali  solution  again 
run  in  to  a  faint  indication,  which  if  permanent  on  the  addition 
of  more  mannite,  may  be  taken  as  the  reading  point.  About 
1-2  grm.  of  mannite  are  necessary  for  a  determination.  The 
boro-mannite  compound  is  sufficiently  acidic  to  liberate  iodine 
abundantly,  but  it  appears  to  be  a  time  reaction,  and  at 
the  end  of  six  hours  only  about  95  per  cent  of  the  theoretical 
amount  (considering  B2O3  as  a  bivalent  acid)  has  been  thrown 
out.  The  combination  of  boric  acid  and  mannite  liberates  in 
the  presence  of  iodide  and  iodate  immediately  only  about 
one-half  the  iodine  required  on  the  theory  that  B2O8  under  these 
conditions  acts  as  a  bivalent  acid,  or  with  the  neutralizing 
power  of  metaboric  acid,  HOBO.  When  no  mannite  is  present 
phenolphthalein  gives  an  alkaline  indication  when  only  about 
one-half  the  amount  of  alkali  theoretically  necessary  to  form 
the  metaborate,  NaOBO,  has  been  added.  Obviously,  then,  the 
starch  iodide  coloration  will  not  appear  at  all  on  the  addition 
of  mannite,  if  about  one-half  the  free  boric  acid  is  first 
neutralized  by  the  solution  of  alkali,  and  the  remainder  of  the 
alkali  immediately  added  to  complete  neutralization.  The 
point  at  which  the  danger  of  the  appearance  of  the  iodide 
coloration  on  the  addition  of  mannite  has  been  passed,  is 
roughly  indicated  before  the  mannite  has  been  added  by  the 
appearance  of  the  strong  alkaline  indication  of  phenolphthalein. 
This  indicator  would  not  need  to  be  added  at  all,  if  the 
boromannite  compound  quickly  and  regularly  liberated  iodine 
from  the  iodide  and  iodate.  The  fact,  however,  that  this 
compound  of  boric  acid  and  mannite  —  as  has  been  ascertained 
by  experiment  —  liberates,  on  standing  twelve  hours,  about  99 
per  cent  of  the  theoretical  amount  of  iodine,  places  the  strength 
of  this  acid  above  that  of  citric  or  tartaric  acid  as  investigated 
by  Furry.*  With  phenolphthalein,  however,  the  end  reaction 
is  sharp  and  the  small  amount  of  carbonate  present  in  the 
standard  solution  of  alkali  is  precipitated  by  the  barium 
*  Am.  Chem.  Jour.,  vi,  341. 


188  A    VOLUMETRIC  METHOD  FOR   THE 

chloride  already  in  the  solution.  The  calculation  must 
therefore  be  based  on  the  amount  of  free  hydroxide  in  the 
standard  solution  of  alkali  used,  according  to  the  following 
representation : 

B208  +  2NaOH  =  2NaOBO  +  H20. 

The  best  results  and  the  most  definite  indications  are 
obtained  in  cold  solution  of  a  volume  not  greater  than  50 
cm3.  This  fact  accords  with  the  observations  of  Magnanini  * 
that  the  relative  electrical  conductivity  of  the  boro-mannite 
solution  is  decreased  by  dilution  and  elevation  of  the 
temperature.  When  silicates  are  present  in  solution,  the 
silicondioxide  is  liberated  by  the  excess  of  hydrochloric  acid, 
and  this  oxide,  whether  in  hydrous  or  anhydrous  condition, 
neither  affects  the  indication  with  iodine  nor  phenolphthalein, 
nor  does  it  form  with  mannite  a  compound  of  acidic  proper- 
ties. Ammonium  salts  interfere  with  the  indication  given 
by  phenolphthalein  and  may  be  removed  by  boiling  with 
potassium  hydroxide  in  excess,  or  an  indicator  used  not 
affected  by  them. 

To  test  the  action  of  fluorides  in  the  process,  several 
experiments  were  made  in  which  hydrofluoric  acid  (10  cm3  of 
T^  solution)  was  introduced  into  the  solution  containing  salts 
of  sodium,  free  hydrochloric  and  boric  acids.  Barium  chloride 
was  then  added  and  the  analysis  for  boric  acid  completed  in 
the  usual  way  without  the  accuracy  of  the  results  being  in  any 
way  interfered  with  by  the  presence  of  hydrofluoric  acid. 

Table  I  contains  the  results  of  a  series  of  analyses  in  which 
the  boric  acid  was  first  drawn  into  a  excess  of  sodium  hy- 
droxide, then  estimated  according  to  the  method  described. 

The  standard  solutions  of  boric  acid  used  contained,  in  A, 
7.153  grm.,  and  in  B,  7.706  grm.  per  liter.  The  solution  of 
free  sodium  hydroxide  was  0.21427  normal. 

Practical  tests  of  the  method  upon  specimens  of  crude 
calcium  borate  and  colemanite  are  recorded  in  Table  II  and 
Table  III. 

*  Gaz.  chim.  ital.,  xx,  428,  and  xxi,  134. 


ESTIMATION  OF  BORIC  ACID. 
TABLE  I. 


189 


Exp. 

B,08  SoL 

taken. 

NaOH  SoL 
required. 

SSL 

B,0, 

found. 

Errors  on 
B20S. 

A 

f(l) 
j  2) 

1(3) 

ClUg 

21.95 
20.68 
20.73 

cms. 
21.02 
19.65 
19.63 

grin* 

0.1671 
0.1479 
0.1483 

grin. 

0.1577 
0.1474 
0.1473 

grm. 
0.0006+ 
0.0005- 
0.0010- 

B 

- 

(4) 
(5) 
(6) 
(7) 
(8) 
,(9) 

23.05 
23.10 
22.76 
24.08 
22.00 
20.78 

23.71 
23.80 
23.35 
24.78 
22.60 
21.28 

0.1776 
0.1780 
0.1754 
0.1855 
0.1695 
'  0.1601 

0.1777 
0.1783 
0.1750 
0.1857 
0.1686 
0.1695 

0.0001+ 
0.0003+ 
0.0004- 
0.0002+ 
0.0009- 
0.0006- 

TABLE  n. 
ANALYSIS  OP  CRUDE  BORATE  OF  LIME. 


Bxp. 

Ca  borate 
taken. 

B2O8  found. 

B20, 

*%• 

grm. 

grm. 

(1) 

0.4016 

0.2289 

56.99 

(2) 

0.4044 

0.2302 

56.92 

(3) 

0.4000 

0.2285 

57.11 

TABLE  m. 
ANALYSIS  OP  COLEMANITE. 


Bxp. 

Mineral  taken. 

B2O8  found. 

%B,08. 

Average. 

grm. 

grm. 

(1) 

0.4034 

0.2064 

61.15] 

(2) 

0.4070 

0.2069 

60.80 

(3 

0.6004 

0.3054 

50.86 

ff\  QQOf 

(4 

0.6006 

0.3056 

50.89 

ou.yy% 

(5 

0.5059 

0.2592 

61.24 

6 

0.5092 

0.2592 

60.89  J 

The  finely-ground  minerals  were  dissolved  in  hydrochloric 
acid  and  the  analyses  proceeded  with  as  described  above. 

An  analysis  for  boric  acid  by  this  process  can  be  completed 
in  five  minutes  and  the  results  are  obviously  accurate  within 
the  limits  of  ordinary  analysis. 

The  usually  interfering  substances,  fluorine,  silica,  and 
carbon  dioxide,  have  no  detrimental  influence  on  the  results 
of  this  process. 


XXIV 

THE  CONSTITUTION  OF  THE  AMMONIUM 
MAGNESIUM  PHOSPHATE  OF  ANALYSIS. 

BY  F.  A.  GOOCH  AND  MARTHA  AUSTIN.* 

IN  a  recent  paper  from  this  laboratory  f  it  has  been  shown 
that  the  presence  of  ammonium  chloride  or  other  ammonium 
salt  is  necessary  in  the  precipitation  of  manganese  as  the 
ammonium  manganese  phosphate  by  microcosmic  salt  in 
order  that  the  precipitate  may  have  the  ideal  constitution 
represented  by  the  symbol  NH4MnPO4. 

It  was  also  shown  that  the  solvent  effect  of  the  ammonium 
chloride  upon  the  precipitated  ammonium  manganese  phosphate 
is  not  marked  when  an  excess  of  the  precipitant  is  present  in 
solution. 

The  relations  disclosed  in  this  paper  suggest  that  the 
chemical  constitution  of  the  precipitate  rather  than  mechanical 
contamination  and  varying  solubility — the  explanations  gene- 
rally accepted,  and,  indeed,  advocated  by  one  of  us  in  a  former 
paper  J  —  may  be  responsible  for  observed  variations  in  the 
weight  of  the  residue  derived  by  the  ignition  of  the  similar 
salt  of  magnesium,  the  ammonium  magnesium  phosphate, 
precipitated  by  an  excess  of  a  soluble  phosphate  from  the 
solution  of  a  magnesium  salt,  or  from  the  solution  of  a  soluble 
phosphate  by  an  excess  of  a  magnesium  salt. 

Precipitation  by  Excess  of  the  Soluble  Phosphate. 

The  precipitation  of  the  magnesium  salt  by  an  excess  of  the 
soluble  phosphate  was  first  studied.  For  this  work  a  solution 
of  pure  magnesium  nitrate  was  prepared  by  dissolving  the 

*  From  Am.  Jour.  Sci.,  vii,  187. 

t  Am.  Jour.  Sci.,  vi,  233.    This  volume,  p.  121. 

J  Am.  Chem.  Jour.,  i,  391. 


AMMONIUM  MAGNESIUM  PHOSPHATE. 


191 


pure  magnesium  oxide  of  commerce  in  a  slight  excess  of  pure 
hydrochloric  acid  and  boiling  with  more  magnesium  oxide. 
After  filtering  off  the  excess  of  magnesium  oxide  and  any  trace 
of  iron  or  members  of  the  higher  groups,  the  solution  was 
precipitated  by  ammonium  carbonate,  the  precipitate  was 
washed  by  repeated  boilings  and  nitrations  until  silver  nitrate 
gave  no  precipitate  in  the  solution  acidified  with  nitric  acid. 
This  precipitated  carbonate  was  nearly  dissolved  in  nitric  acid 
and  the  solution  was  boiled  with  an  excess  of  the  carbonate 
(for  the  purpose  of  removing  traces  of  barium,  strontium, 
and  calcium)  filtered,  and  diluted  to  definite  volume.  The 
evaporation  of  a  definite  volume  of  the  solution  and  strong 
ignition  of  the  residue  would  be  a  most  natural  method  of 
establishing  a  standard  of  the  solution,  were  it  not  for  the 
fact,  pointed  out  by  Richards  and  Rogers,*  that  the  oxide  of 
magnesium  retains  on  ignition  occluded  nitrogen  and  oxygen 
enough  to  increase  its  weight  sensibly.  For  this  reason  the 
nitrate  was  converted  to  the  sulphate  and  weighed  as  such  — 
either  by  evaporating  to  dryness  in  a  weighed  platinum 
crucible  a  definite  volume  of  the  solution,  igniting  as  oxide, 
and  changing  to  the  sulphate  by  heating  with  sulphuric  acid ; 
or,  by  evaporating  the  magnesium  nitrate  directly  with  an 
excess  of  sulphuric  acid  of  half  strength.  In  this  treatment 
the  excess  of  acid  was  removed  by  heating  the  platinum 
crucible  upon  a  porcelain  ring  or  triangle  so  placed  within  a 
porcelain  crucible  that  the  bottom  and  walls  of  the  inner 
crucible  were  distant  about  one  centimeter  from  the  bottom 

TABLE  I. 


MgS04  obtained  by 
converting  ignited 
MgO  into  the  sulphate. 

MgSO4  obtained 
directly  from  50  cm3 
Mg(N03)s. 

Theoretical 
amount  of  MgO 
in  MgS04. 

grm. 
0.5748 
0.6739 

grm. 

0.5741 
0.5750 

grin. 

0.1924 
0.1923 
0.1922 
0.1925 

*  Amer.  Chem.  Jour.,  xvi,  567. 


192  CONSTITUTION  OF  THE  AMMONIUM 

and  walls  of  the  outer  crucible.  The  excess  of  acid  is  easily 
removed  in  this  way,  and  the  outer  crucible  may  be  heated 
to  redness  without  danger  of  breaking  up  the  magnesium 
sulphate.  The  results  of  this  work,  taking  O  =  16,  Mg  =  24.3, 
N  =  14.03,  S  =  32.06,  are  given  in  the  accompanying  table. 

The  magnesium  oxide  obtained  by  direct  ignition  of  the 
nitrate  weighed  on  the  average  about  0.0010  grm.  more  than 
the  oxide  theoretically  present  in  the  weighed  sulphate  from 
equal  portions  of  the  solution. 

Before  proceeding  to  study  possible  chemical  effects  of 
ammonium  chloride  in  determining  the  constitution  of  the 
ammonium  magnesium  phosphate,  it  is  obviously  necessary  to 
define  the  extent  to  which  the  ammonium  salt  may  exert  a 
solvent  action  in  presence  of  the  precipitant.  Fresenius 
estimated  that  ammonium  magnesium  phosphate  is  soluble  in 
15293  parts  of  cold  water,  but  the  method  of  investigation 
employed  did  not  entirely  preclude  the  possibility  of  counting 
as  ammonium  magnesium  phosphate  soluble  material  included 
and  held  in  the  original  precipitate.*  According  to  Kissel  f 
the  phosphate,  which  dissolves  in  a  mixture  of  ammonia  and 
water  in  the  proportion  of  0.0040  grams  to  the  liter  and  in  the 
proportion  of  0.0110  grams  to  the  liter  in  a  similar  mixture 
containing  also  18  grm.  of  ammonium  chloride,  is  practically 
insoluble  in  the  latter  mixture  if  an  excess  of  magnesia 
mixture  be  added;  and  HeintzJ  showed  that  the  effect  of 
adding  an  excess  of  sodium  phosphate  in  the  solution  is 
similar. 

So  far  as  appears,  no  quantitative  experiments  have  been 
recorded  hi  which  the  behavior  of  a  mixture  of  ammonium 
chloride  and  magnesium  salt  and  an  insoluble  phosphate  in  a 
solution  only  slightly  ammoniacal  has  been  tested,  though  as  a 
matter  of  convenience  the  use  of  faintly  ammoniacal  solutions 
and  faintly  ammoniacal  washwater  is  to  be  preferred  to  the 
mixture  of  strong  ammonia  and  water  [1  :  3]  ordinarily 
employed.  As  a  preliminary  step,  therefore,  in  the  work  to  be 

*  Fresenius,  6te  Aufl.,  p.  805.  t  Zeitschr.  anal.  Chem.,  viii,  173. 

J  Zeitschr.  anal.  Chem.,  ix,  16. 


MAGNESIUM  PHOSPHATE   OF  ANALYSIS.        193 


described,  experiments  were  made  to  find  how  small  an  amount 
of  magnesium  could  be  detected  in  solution  by  precipitating 
with  microcosmic  salt,  either  alone  or  in  presence  of  ammonium 
chloride  in  faintly  ammoniacal  solutions.  The  ammonium 
chloride  used  for  these  tests  (as  well  as  in  the  similar  quanti- 
tative work  following)  was  purified  by  boiling  with  a  faint 
excess  of  ammonia,  filtering,  digesting  twelve  hours  with 
microcosmic  salt,  and  filtering  again.  The  results  are  given 
in  Table  II. 

TABLE  II. 


Weight  of 
MgO  taken  as 
the  nitrate. 

H(NH4)NaP04. 
4H,0  taken. 

Volume. 

NH4C1 
taken. 

Opalescent 
precipitation. 

grm. 

grm. 

cm8 

grm. 

(  0.0003 

1.75 

100 

Marked. 

1  0.0003 

1.75 

500 

Marked. 

(0.0003 

1.75 

100 

10 

Marked. 

}  0.0003 

1.75 

500 

10 

Marked. 

(  0.0003 

1.75 

500 

30 

Faint. 

0.0001 

1.75 

100 

. 

Marked. 

(  0.0001 

1.75 

100 

10 

Marked. 

}  0.0001 

1.75 

500 

10 

Faint. 

(  0.0001 

1.75 

500 

60 

Faint. 

The  results  of  these  tests  show  that  even  so  little  as  0.0001 
grm.  of  magnesium  oxide  may  be  detected  in  five  hundred 
cubic  centimeters  of  faintly  ammoniacal  water  containing  as 
much  as  sixty  grams  of  ammonium  chloride,  f  It  is  plain  that 
strongly  ammoniacal  liquids  are  entirely  unnecessary  in  the 
precipitation  of  the  ammonium  magnesium  phosphate  under 
the  conditions.  In  nearly  all  the  experiments  to  be  detailed 
use  was  made,  therefore,  of  faintly  ammoniacal  solutions  and 
wash-water. 

In  Table  III  are  given  the  results  obtained  in  a  study  of  the 
effects  of  varying  proportions  of  ammonium  chloride  and  the 
soluble  phosphate  upon  the  constitution  of  the  precipitate. 
All  precipitates  were  gathered  upon  asbestos  in  the  filtering 

t  It  was  found  also,  incidentally,  that  the  presence  of  reasonable  amounts  of 
ammonium  oxalate  (100  cm8  of  the  saturated  solution)  does  not  interfere  with 
the  precipitation  of  the  ammonium  magnesium  phosphate  by  microcosmic 
salt. 

VOL.   XI.  —  13 


194 


CONSTITUTION  OF  THE  AMMONIUM 


crucible,  washed  in  faintly  ammoniacal  water,  and  ignited  as 
usual.  In  every  case  the  precipitation  was  practically  com- 
plete ;  for,  upon  allowing  the  filtrates  with  the  wash-water  to 
stand  for  several  days  after  further  addition  of  microcosmic 
salt,  nothing  but  insignificant  traces  of  a  precipitate  —  not 
exceeding  0.0001  grm.  —  ever  appeared.  In  the  experiments 
of  section  A  precipitations  were  made  in  the  cold  by  the  action 
of  microcosmic  salt  in  considerable  excess  upon  the  solutions 
of  magnesium  nitrate  containing  varying  amounts  of  ammonium 
chloride.  In  experiments  (1)  to  (5)  the  liquid  was  made 
faintly  ammoniacal  after  the  addition  of  the  precipitant  and 
the  precipitate  was  filtered  off  immediately  after  complete 
subsidence;  in  experiments  (6)  to  (10)  the  precipitate  first 

TABLE  in. 


Exp. 

Mg2P207 
corresponding 
to  Mg(N03)a 

Mg2P207 
found. 

Error  in 
terms  of 
Mg2P207. 

Error  in 
terms  of 
MgO. 

NH4C1 
present. 

HNaNH4P04. 
4H2O  used. 

Volume. 

A. 

grm. 

grm. 

grin. 

grm. 

grm. 

grm. 

cm3. 

(1) 

0.5311 

0.5418 

0.0107+ 

0.0038+ 



2.5 

150 

2) 

0.5311 

0.5462 

0.0151+ 

0.0057+ 

2 

2.5 

150 

8 

0.5311 

0.5408 

0.0097+ 

0.0035+ 

2 

2.5 

150 

(4) 

0.5311 

0.5500 

0.0189+ 

0.0068+ 

60 

2.5 

250 

(5) 

0.5311 

0.5520 

0.0209+ 

0.0075+ 

60 

2.5 

250 

(6) 

0.5311 

0.5345 

0.0034+ 

0.0012+ 

2.5 

150 

(7 

0.5311 

0.5371 

0.0060  f 

0.0022+ 

2.5 

150 

(8) 

0.5311 

0.5384 

0.0073+ 

0.0026+ 

2.5 

150 

(9) 

0.5311 

0.5386 

0.0075+ 

0.0027+ 

2.5 

150 

(10) 

0.5311 

0.5415 

0.0104+ 

0.0037+ 

2.5 

150 

B. 

(11) 

0.5311 

0.5312 

0.0001+ 

0.0000 

#  

2.5 

150,100 

(12) 

0.5311 

0.5311 

0.0000 

0.0000 

#  

2.5 

150,100 

(13) 

0.5311 

0.5346 

0.0035+ 

0.0013+ 

2  +  2 

2.5 

150,100 

(14) 

0.5311 

0.5348 

0.0037+ 

0.0014+ 

2+  2 

2.5 

150,100 

(15) 

0.5311 

0.5383 

0.0072+ 

0.0026+ 

5+  6 

2.5 

150,100 

(16) 

0.5311 

0.5368 

0.0057+ 

0.0021+ 

5+  5 

2.5 

150,100 

(17) 

0.5311 

0.5376 

0.0065+ 

0.0023+ 

10  +  10 

2.5 

200,100 

(18) 

0.5311 

0.5395 

0.0084+ 

0.0030+ 

10  +  10 

2.5 

200,100 

(19) 

0.5311 

0.5396 

0.0085+ 

0.0031+ 

60+  5 

2.5 

250,100 

(20) 

0.5311 

0.5389 

0.0078+ 

0.0028+ 

60+  6 

2.6 

250,100 

*  Probably  less  than  1  grm. 


MAGNESIUM  PHOSPHATE  OF  ANALYSIS.       195 

thrown  down  was  redissolved  in  a  very  little  hydrochloric 
acid  and  reprecipitated  by  dilute  ammonia  (the  operation 
being  repeated  several  times)  with  a  view  to  improving  the 
crystalline  condition  of  the  precipitate,  and  this  treatment 
introduced,  of  course,  a  small  amount  of  ammonium  chloride, 
probably  less  than  a  gram.  It  will  be  observed  that  errors  of 
excess  appear  in  all  of  these  determinations,  those  being  the 
greatest  in  the  experiments  in  which  the  largest  amounts  of 
the  ammonium  salt  were  present. 

In  the  experiments  of  section  B  the  manipulation  was  so 
changed  that  the  supernatant  liquid  was  poured  off  (through 
the  filtering  crucible  which  was  to  be  used  subsequently  to 
collect  the  phosphate)  after  the  precipitate  had  subsided  and 
the  insoluble  phosphate  was  dissolved  in  hydrochloric  acid  and 
brought  down  again,  after  dilution,  by  the  addition  of  a  faint 
excess  of  dilute  ammonia.  By  thus  removing  the  supernatant 
liquid  after  the  first  precipitation,  the  excess  of  the  precipitant 
and  the  amounts  of  ammonium  chloride  orginally  present  were 
reduced  to  relatively  low  limits,  so  that  their  effects  in  the 
reprecipitation  were  at  a  minimum,  and  by  adding  varying 
amounts  of  ammonium  chloride,  or  none  at  all,  before  the 
reprecipitation,  it  became  possible  to  demonstrate  the  individual 
effect  of  that  reagent  apart  from  that  of  an  excess  of  the 
microcosmic  salt.  It  will  be  noted  that  in  experiments  (11) 
and  (12),  in  which  no  ammonium  salt  was  added  after  the 
decantation  from  the  first  precipitate,  the  results  are  ideal,  and 
that  the  errors  of  excess  advance  as  the  amounts  of  ammonium 
salt  present  in  the  final  precipitation  increase.  The  quantity 
of  the  ammonium  salt  present  during  the  first  precipitation 
does  not  influence  the  error  in  the  final  precipitation  unless  it 
is  so  large  that  a  simple  decantation  of  the  supernatant  liquid 
would  naturally  leave  an  appreciable  amount  of  it  to  act  when 
the  second  precipitation  takes  place. 

It  is  plain  that  the  errors  of  excess  which  appear  when 
either  the  ammonium  chloride  or  the  soluble  phosphate  is 
present  in  considerable  amount,  must  be  due  either  to  mechan- 
ical inclusion  on  the  part  of  the  highly  crystalline  precipitate, 


196  CONSTITUTION  OF  THE  AMMONIUM 

or  to  variation  in  the  ammonium  magnesium  phosphate  from 
the  ideal  constitution  toward  a  condition  represented  by  a 
phosphate  richer  in  ammonia  and  correspondingly  deficient  in 
magnesium.  If  any  appreciable  amount  of  the  ammonium 
chloride  present  were  held  by  the  precipitate,  it  would  natu- 
rally be  represented  by  magnesium  chloride  after  ignition,  but, 
in  no  one  of  these  experiments,  even  in  those  dealing  with 
sixty  grams  of  ammonium  chloride,  did  the  residue,  after  dis- 
solving in  nitric  acid,  give  with  silver  nitrate  evidence  of  the 
presence  of  more  than  a  mere  unweighable  trace  of  chloride. 
A  special  experiment,  moreover,  in  which  an  attempt  was 
made  to  determine  the  silver  chloride  precipitated  from  the 
solution  hi  nitric  acid  of  an  unignited  precipitate  thrown  down 
by  microcosmic  salt  in  presence  of  sixty  grams  of  ammonium 
chloride,  confirms  this  conclusion :  'the  precipitate  was  unweigh- 
able. If  ammonium  chloride  present  in  the  solution  to  so 
great  an  amount  is  not  included  in  the  precipitate  in  signifi- 
cant quantity,  it  would  seem  to  be  unnatural  that  the  micro- 
cosmic  salt  should  be  included  mechanically  in  any  very  great 
amount.  But  unless  the  microcosmic  salt  was  mechanically 
included,  the  increase  in  weight  must  be  due  to  the  chemical 
influence  of  the  reagents  —  that  is,  to  the  production  of  a 
phosphate  rich  in  ammonium  and  deficient  in  magnesium. 
Berzelius*  recognized  the  existence  of  such  a  phosphate  of 
magnesium ;  but  Wachf  in  following  the  work  of  Berzelius, 
failed  to  find  it.  It  would  be  natural  to  expect  its  formation, 
if  ever,  when  the  precipitating  phosphate  is  in  excess  and 
ammonium  salts  are  present  in  abundance,  with  free  ammonia. 
Obviously  the  natural  effects  of  all  these  reagents  would  be 
toward  the  production  of  a  salt  holding  more  ammonia  and 
more  phosphoric  pentoxide  for  a  given  amount  of  magnesium. 
The  results  of  the  table  seem  to  point  strongly  to  such  ten- 
dencies, and,  by  inference,  toward  the  existence  of  such  a 
compound.  Thus  in  experiments  (11)  and  (12),  in  which  the 

*  Berzelius,  Jahresbericht,  3.  Jahrgang  (1824),  iibersetzt  von  C.  G.  Gmelin, 
8.92. 

t  Schweigger,  1830,  Band  29,  s.  265. 


MAGNESIUM  PHOSPHATE  OF  ANALYSIS.       197 

greater  part  of  this  excess  of  microcosmic  salt  was  removed 
by  decantation  before  the  second  precipitation,  while  no  am- 
monium chloride  was  present  excepting  the  small  amount 
made  by  the  solution  and  reprecipitation  of  the  first  precipi- 
tate, the  error  is  practically  nothing.  In  experiments  (13)  and 
(14),  (15)  and  (16),  (17)  and  (18),  all  similar  to  (11)  and 
(12)  excepting  that  ammonium  chloride  was  present,  the  aver- 
age errors  (+0.0036  grm.  in  terms  of  magnesium  phosphate, 
+0.0064  grm.,  +0.0074,  respectively)  increase  as  the  ammonium 
chloride  is  increased  in  the  final  precipitation.  In  experiments 
(19)  and  (20),  in  which  the  ammonium  chloride  amounted  to 
sixty  grams  in  the  first  precipitation  and  to  five  grams  in  the 
second  in  addition  to  the  amount  that  would  naturally  remain 
after  decanting  the  strong  solution  of  the  former  precipita- 
tion, the  similarity  of  this  error  (+0.0082  in  the  mean)  to 
that  of  the  experiments  in  which  smaller  amounts  of  the 
ammonium  chloride  were  used  throughout  goes  to  show  that 
only  the  amount  of  ammonium  salt  present  in  the  final  pre- 
cipitation counts.  Further,  a  comparison  of  corresponding 
experiments  of  A  and  B  shows  very  plainly  that  the  treatment 
which  involves  the  removal  of  the  large  part  of  the  micro- 
cosmic  salt,  the  solution  of  the  precipitate,  and  reprecipitation 
tends  to  reduce  the  higher  indications.  Thus,  for  example  the 
error  in  (2)  and  (3)  is  +0.0124  gram  in  terms  of  magnesium 
pyrophosphate,  while  in  (13)  and  (14),  similarly  carried  out 
except  the  decantation  of  the  excess  of  the  precipitant,  solu- 
tion and  reprecipitation,  the  error  is  +0.0036  grm. 

The  special  influence  of  free  ammonia  during  precipitation, 
was  investigated  in  the  following  experiments.  Definite 
volumes  of  the  magnesium  nitrate  solution  were  drawn  from  a 
burette  into  a  platinum  dish,  ammonium  chloride  — 10  grm.  — 
was  added,  the  magnesium  was  brought  down  by  dilute 
ammonia  in  presence  of  microcosmic  salt,  and  strong  ammonia 
equal  to  one-third  the  volume  of  the  solution  was  added.  The 
solutions,  after  standing,  were  filtered  off  on  asbestos  under 
pressure  in  a  perforated  crucible,  and  the  precipitates  were 
washed  with  ammonia  diluted  to  the  proportion  of  three  parts 


198 


CONSTITUTION  OF  THE  AMMONIUM 


of  water  to  one  of  ammonia,  dried  after  moistening  with  a 
drop  of  saturated  solution  of  ammonium  nitrate,  ignited  and 
weighed.  The  results  are  given  in  experiments  (1)  and  (2)  of 
Table  IV.  In  these  determinations  the  mean  error  reaches 
+0.0193  gnn.  in  terms  of  magnesium  pyrophosphate ;  while 
in  experiments  (3)  and  (4),  made  similarly  excepting  that  the 
supernatant  liquid  was  decanted  from  the  precipitate  first 
thrown  down,  the  precipitate  dissolved  in  hydrochloric  acid, 
and  after  dilution  reprecipitated  by  dilute  ammonia  imme- 
diately supplemented  by  enough  strong  ammonia  to  make  one- 
fourth  the  volume  of  the  entire  solution,  the  error  amounts  in 
the  mean  to  +0.0061  in  terms  of  the  pyrophosphate. 

TABLE  IV. 


Exp. 

Mg2P207 
corresponding 
to  Mg(NO3)a 

Mg2P207 
found. 

Error  in 
terms  of 
Mg2P207. 

Error  in 
terms 
of  MgO. 

NH4C1 
present. 

HNaNH4P04. 
4H20  used. 

Volume. 

(1) 

gnn. 

0.5311 

grm. 
0.5503 

grm. 
0.0192+ 

gnn. 
0.0069+ 

grm. 

10 

grm. 
2.5 

cms 
200 

(2) 

0.5311 

0.5505 

0.0194+ 

0.0070+ 

10 

2.5 

200 

(3) 

0.6311 

0.5393 

0.0082+ 

0.0029+ 

10,— 

2.5 

200,100 

(4) 

0.5311 

0.5351 

0.0040+ 

0.0017+ 

10,- 

2.5 

200,100 

In  experiments  (1)  and  (2)  the  precipitate  was  influenced  by 
an  excess  of  microcosmic  salt,  ammonium  chloride,  and  free 
ammonia  in  large  amount;  in  experiments  (3)  and  (4),  by 
decanting  in  the  manner  previously  described,  by  dissolving 
the  precipitate,  and  reprecipitating,  the  effects  of  an  excess  of 
microcosmic  salt  and  ammonium  chloride  are  reduced  to  a  mini- 
mum, and,  in  a  comparison  of  the  results  with  those  of  experi- 
ments (11)  and  (12)  of  Table  III  the  tendency  of  the  free 
ammonia  comes  to  view.  The  results  discussed  seem  certainly 
to  point  to  a  general  tendency  on  the  part  of  free  ammonia, 
ammonium  choride  and  excess  of  the  phosphate  to  produce  a 
salt  rich  in  ammonia  and  deficient  in  magnesium,  which  for  a 
definite  amount  of  magnesia  precipitated  must  leave  upon  igni- 
tion a  residue  weighing  more  than  the  normal  phosphate. 
If  it  be  assumed  that  a  salt  of  the  symbol  (NH4)4Mg(PO4)2 


MAGNESIUM  PHOSPHATE  OF  ANALYSIS.        199 

—  the  next  natural  step  to  the  normal  salt,  NH4MgPO4  — 
is  present  in  the  precipitate,  the  residue  which  such  a  salt 
would  leave  upon  ignition  would  be  the  metaphosphate 
Mg(PO3)2.  From  the  relations  of  the  symbols  for  magnesium 
pyrophosphate  and  magnesium  metaphosphate  the  weight  of 
the  residue  obtained,  and  the  weight  of  the  pyrophosphate 
theoretically  derivable  from  the  weight  of  magnesium  salt 
used,  it  is  possible,  of  course,  to  calculate  the  proportionate 
amounts  of  pyrophosphate  and  metaphosphate  present  in  any 
ignited  residue.  Proceeding  in  this  manner,  it  appears  that, 
in  order  to  account  for  the  variations  noted,  it  is  necessary 
to  assume  the  presence  in  many  cases  of  very  considerable 
amounts  of  the  metaphosphate.  Thus,  in  the  case  of  those 
results  obtained  according  to  the  usually  accepted  method  of 
precipitating  and  washing  with  strongly  ammoniacal  liquids, 
viz.,  in  experiments  (1)  and  (2)  of  Table  IV,  the  proportion 
of  metaphosphate  needed  to  account  for  the  observed  error 
reaches  ten  per  cent. 

Precipitation  ~by  Excess  of  the  Magnesium  Salt. 

The  relations  which  obtain  in  the  reverse  process  of  pre- 
cipitation—  the  action  of  an  excess  of  the  magnesium  salt 
upon  a  soluble  phosphate  —  were  studied  in  experiments  to 
be  described.  A  solution  of  pure  hydrogen  disodium  phos- 
phate was  prepared  by  carefully  recrystallizing  the  pure  salt 
of  commerce  five  tunes  from  distilled  water  in  a  platinum 
dish,  dissolving  the  crystals,  and  diluting  to  definite  volume. 
The  standard  of  the  solution  was  established  by  evaporating 
to  dryness  in  a  weighed  platinum  crucible  known  volumes 
of  the  solution,  igniting  the  residue  and  weighing  the  sodium 
pyrophosphate.  Magnesia  mixture,  the  precipitant,  was  pre- 
pared by  dissolving  fifty-five  grams  of  magnesium  chloride 
in  as  little  water  as  possible  and  filtering,  mixing  with  this 
solution  twenty-eight  grams  of  ammonium  chloride  purified 
by  treating  it  in  strong  solution  with  bromine  water  and  a 
slight  excess  of  ammonia,  filtering,  diluting  to  one  liter,  and, 
after  standing  for  some  hours,  filtering  again. 


200 


CONSTITUTION  OF  THE  AMMONIUM 


The  tests  of  the  following  table  show  that  the  precipitation 
of  a  soluble  phosphate  by  the  magnesia  mixture  is  practically 
complete  in  faintly  ammoniacal  solutions  even  when  very 
dilute  and  charged  with  large  amounts  of  ammonium  chloride, 
provided  the  magnesia  mixture  is  present  in  sufficiently  large 
excess. 

TABLE  V. 


P2O,  in  HNa,PO4 
taken. 

Magnesia 
mixture. 

Volume. 

NHtCl. 

Precipitation  visible. 

grm. 

cms 

cm8 

grm. 

0.0005 

10 

100 

t 

At  once  throughout 

0.0005 

50 

100 

. 

the  liquid. 

0.0005 

10 

100 

10 

it 

0.0005 

10 

200 

60 

tt 

0.0001 

50 

250 

60 

u 

0.0001 

10 

100 

« 

0.0001 

10 

100 

10 

tl 

0.0001 

50 

200 

10 

« 

0.0001 

10 

250 

60 

« 

0.0001 
0.0001 

50 
50 

300 
500 

60 
60 

After  settling  out. 

This  conclusion  was  further  substantiated  by  an  actual  test 
(by  the  molybdate  method)  of  the  ignited  residue,  obtained 
by  evaporating  a  filtrate  from  ammonium  magnesium  phos- 
phate (equivalent  to  0.8614  grm.  of  the  pyrophosphate) 
precipitated  by  a  faintly  ammoniacal  solution  of  magnesia 
mixture  in  presence  of  60  grm.  of  ammonium  chloride, 
which  gave  a  precipitate  of  ammonium  phosphomolybdate 
yielding  0.0002  grm.  of  magnesium  pyrophosphate.  It  is 
evident,  therefore,  that  any  considerable  deficiencies  of  weight 
of  the  magnesium  phosphate  obtained  by  precipitating  equal 
amounts  of  a  soluble  phosphate  by  magnesia  mixture  in 
presence  of  varying  amounts  of  ammonium  chloride,  cannot 
be  attributed  to  varying  solubility  of  the  magnesium  phosphate 
under  changing  proportions  of  the  ammonium  chloride. 

The  results  recorded  in  section  A  of  Table  VI  were  obtained 
by  treating  definite  volumes  of  the  pure  solution  of  hydrogen 
disodium  phosphate  with  magnesia  mixture,  in  slight  excess 
above  the  amount  required  to  bring  down  the  phosphate,  and 


MAGNESIUM  PHOSPHATE  OF  ANALYSIS. 


201 


making  the  solution  distinctly  ammoniacal.  After  thorough 
subsidence,  the  precipitate  was  filtered  off  on  asbestos  under 
pressure  in  a  perforated  platinum  crucible,  washed  hi  water 
faintly  ammoniacal,  dried,  ignited  and  weighed.  In  experi- 
ments (1),  (5)  and  (6),  only  the  ammonium  chloride  present  in 
the  magnesia  mixture  was  used ;  in  the  other  cases  weighed 
portions  were  added.  In  the  experiments  of  section  B,  the 
precipitate  was  dissolved  in  hydrochloric  acid  after  filtering  off 
the  supernatant  liquid,  brought  down  again  in  dilute  solution 
by  ammonia  in  distinct  excess,  and  thereafter  treated  as  in  the 
experiments  of  section  A.  The  experiments  of  section  C 
were  conducted  similarly  to  (1),  (5)  and  (6)  of  A  excepting 
that  the  magnesium  mixture  was  introduced  into  the  am- 

TABLE  VL 


Exp. 

Mg2P207 
corre- 
sponding 
to 

taken. 

found. 

Error  in 
terms  of 
MglP,07. 

Error  in 
terms  of  P. 

Volume. 

NH4C1 

in  mag- 
nesia 
mix- 
ture. 

NH^l 
added. 

MgCl, 
6H2O 
in  mag- 
nesia 
mix- 
ture. 

A 

(1) 
(2) 
(3) 
(4) 
(5) 
(6) 
7) 
8) 
9) 

grm. 
0.8615 
0.8615 
0.8615 
0.8615 

0.0852 
0.0852 
0.0852 
0.0852 
0.0852 

gnu. 
0.8613 
0.8615 
0.8602 
0.8561 

0.0862 
0.0866 
0.0847 
0.0830 
0.0811 

grm. 

0.0002- 
0.0000 
0.0013- 
0.0054- 

0.0010+ 
0.0014+ 
0.0005- 
0.0022- 
0.0041- 

grm. 

0.00005- 
0.00000 
0.00036- 
0.00151- 

0.00028+ 
0.00039+ 
0.00014- 
0.00062- 
0.00115- 

cm3 
150 
200 
200 
300 

100 
100 
200 
200 
300 

griii. 

1.68 
1.68 
1.68 
1.68 

0.28 
0.28 
0.28 
0.28 
0.28 

griii. 

*20' 
20 
60 

*20' 
20 
60 

grm. 
3.3 
3.3 
3.3 
3.3 

0.55 
0.55 
0.55 
0.55 
0.55 

B 

(10) 

(11) 
(12) 
(13) 

(14) 
(15) 
(16) 
(17) 

0.8111 
0.8615 
0.8615 
0.8615 

0.0852 
0.0852 
0.0852 
0.0852 

0.8114 
0.8613 
0.8578 

0.8487 

0.0855 
0.0656 
0.0853 
0.0819 

0.0003+ 
0.0002- 
0.0037- 
0.0128- 

0.0003+ 
0.0004+ 
0.0001+ 
0.0033- 

0.00008+ 
0.00006- 
0.00103- 
0.00358- 

0.00008+ 
0.00011+ 
0.00003+ 
0.00092- 

150,100 
150,000 
200,100 
200,100 

100,100 
100,100 
150,100 
200,100 

1.68 
1.68 
1.68 
1.68 

0.28 
0.28 
0.28 
0.28 

!  ,20 
.,60 

io]  ! 

10,10 
20,20 

3.3 
3.3 
3.3 
3.3 

0.55 
0.55 
0.55 
0.55 

C 

(18) 
(19) 

0.8111 
0.8111 

0.8071 
0.8052 

0.0040- 
0.0059- 

0.00112- 
0.00165- 

120 
120 

1.4 

1.4 

•  •  • 

2.75 
2.75 

202  CONSTITUTION  OF  THE  AMMONIUM 

moniacal  solution  of  the  phosphate  drop  by  drop  from  a 
burette.  The  precipitations  in  A,  B,  and  C  were  proved  to 
be  practically  complete  ;  for  by  treatment  of  the  filtrates  with 
more  magnesia  mixture  and  standing,  no  more  than  a  trace 
—  0.0001  grm.  at  the  most  —  of  the  phosphate  was  found. 
The  ignited  residues  never  contained  more  than  a  mere  trace 
of  chlorine. 

While  the  results  are  not  entirely  regular,  the  tendency  of 
the  ammonium  salt  to  produce  errors  of  deficiency  in  propor- 
tion to  its  amount  is  plain  if  we  compare  among  themselves 
the  experiments  of  A  upon  similar  amounts  of  phosphate,  and 
then  those  of  B  upon  similar  amounts  of  phosphate  among 
themselves ;  and  by  a  comparison  of  corresponding  results  in 
A  and  B  it  is  clearly  shown  that  the  presence  of  an  excess  of 
magnesia  mixture  tends  to  counteract  more  or  less  completely 
errors  of  deficiency  due  to  the  action  of  the  ammonium  chloride. 
These  facts  are  quite  in  harmony  with  the  hypothesis  that  the 
ammonium  salt  tends  to  produce  an  ammonium  magnesium 
phosphate  richer  in  ammonia  and  phosphoric  acid  and  poorer 
in  magnesia  than  the  normal  salt  NH4MgPO4;  for,  though 
the  production  of  such  a  salt  in  presence  of  an  excess  of  the 
soluble  phosphate  compels  the  combination  of  a  definite 
amount  of  magnesium  with  more  than  the  normal  amounts  of 
phosphoric  acid  and  ammonia  (as  was  the  case  in  the  former 
series  of  experiments),  when  the  supply  of  the  soluble 
phosphate  is  limited  the  amount  of  magnesium  associated  with 
it  must  fall  below  the  normal  (as  is  the  case  in  the  present 
series  of  experiments).  Moreover,  the  behavior  of  the 
precipitant  is  quite  in  accord  with  the  hypothesis ;  for,  though 
the  influence  of  an  excess  of  the  soluble  phosphate  would 
naturally  tend  (as  was  observed)  in  the  same  direction  as  that 
of  the  ammonium  salt  and  free  ammonia,  viz.,  to  the  production 
of  the  phosphate  deficient  in  magnesium,  the  tendency  of  an 
excess  of  the  magnesium  salt  must  obviously  be  to  increase 
the  amount  of  magnesium  in  the  phosphate,  as  was  observed 
in  the  experiments  of  Table  VI.  The  hypothesis  fits  the  facts, 
therefore,  on  both  sides ;  and,  if  precipitation  is  practically 


MAGNESIUM  PHOSPHATE  OF  ANALYSIS.          203 

complete  (as  was  shown  to  be  the  case  throughout)  the 
argument  for  the  existence  of  an  ammonium  magnesium 
phosphate  —  poorer  than  the  normal  salt  in  magnesium  — 
possibly  the  salt  (NH^Mg^O^a  —  seems  to  be  strong. 


The  Practical  Determination   of  Magnesium  and  Phosphoric 

Acid. 

In  determining  magnesium  by  the  procedure  in  ordinary  use, 
the  tendency  is«strong  —  as  is  shown  in  experiments  (1)  and(2) 
of  Table  IV  —  toward  high  plus  errors,  and  the  error  is  due 
to  the  combined  effects  of  excesses  of  the  precipitant,  the 
ammonium  salt,  and  free  ammonia.  The  experiments  (11) 
and  (12)  of  B,  Table  III,  show  conclusively  that  such 
tendencies  to  error  may  be  counteracted  effectively  by  pouring 
off  the  supernatant  liquid  (through  the  filter  to  be  used 
subsequently  to  collect  the  precipitate)  as  soon  as  the  precipi- 
tate subsides,  dissolving  the  phosphate  in  the  least  amount 
of  hydrochloric  acid,  bringing  it  down  again,  after  dilution, 
by  a  faint  excess  of  ammonia,  filtering  (best,  we  think,  on 
asbestos,  under  pressure),  washing  with  faintly  ammoniacal 
water,  and  igniting  as  usual. 

Many  years  ago*  a  method  of  precipitating  the  ammonium 
magnesium  phosphate  was  advocated  by  Professor  Wolcott 
Gibbs,  which  consists,  essentially,  in  boiling  the  solution  of  the 
magnesium  salt  with  microcosmic  salt  and  adding  ammonia 
after  cooling,  and  by  which  most  exact  analytical  results 
were  obtained.  Our  experience  confirms  completely  that  of 
Gibbs,  and  we  desire  to  direct  attention  again  to  a  procedure 
the  advantage  of  which  has,  unfortunately,  not  been  broadly 
known  and  accepted.  Even  in  the  presence  of  considerable 
amounts  of  ammonium  chloride  this  process  yields  a  phosphate 
of  nearly  ideal  constitution  if  only  the  boiling  be  prolonged 
from  three  to  five  minutes.  The  greater  part  of  the  ammonium 
magnesium  phosphate  —  about  90  per  cent  —  forms  in  this 
process  before  free  ammonia  is  added,  and  the  ammonium 
which  enters  the  phosphate  thus  formed  is  derived  from  the 
*  Am.  Jour.  Sci.  [3],  v,  114. 


204  CONSTITUTION  OF  THE  AMMONIUM 

microcosmic  salt,  which  must  become  correspondingly  acidic. 
Under  these  conditions,  the  tendency  to  form  an  insoluble 
ammonium  magnesium  phosphate  richer  in  ammonia  and  poorer 
in  magnesia  than  the  normal  salt,  does  not  develop.  In  the 
process  of  Gibbs,  as  well  as  in  the  modified  precipitation 
process  in  the  cold,  the  use  of  the  faintly  ammoniacal  solution 
and  wash-water  is  sufficient  and  advantageous. 

In  the  precipitation  of  a  soluble  phosphate  by  magnesia 
mixture  the  tendency  of  the  precipitant  and  that  of  the 
ammonium  salt  are  antagonistic,  so  that  the  effect  of  the  latter 
salt  is  somewhat  masked,  though  manifest.  This  opposition 
of  effects  has  been  noted  by  Mahon,*  who,  though  regarding 
the  actual  attainment  of  an  exact  balance  as  uncertain,  ventures 
the  opinion  that  accurate  results  should  be  attainable  by  the 
careful  relative  adjustment  of  the  proportions  of  the  precipitant 
and  ammonium  salt.  Mahon  claims  to  get  the  best  results  by 
a  very  gradual  addition  of  magnesia  mixture  to  the  ammoniacal 
solution  of  the  phosphate  containing  about  sixteen  per  cent  of 
ammonium  chloride,  strong  ammonia  being  added  subsequently. 
From  our  observations,  however,  recorded  in  section  C  of  Table 
VI,  it  appears  that  the  method  of  introducing  the  magnesia 
mixture  gradually  into  the  ammoniacal  phosphate  (taken  in 
quantity  sufficiently  large  to  give  unmistakable  indications) 
produces  a  precipitate  deficient  in  magnesium  and  so  leads  to 
errors  of  deficiency  in  the  phosphorus  indicated.  The  use  of 
strong  ammonia,  moreover,  we  have  shown  to  be  both  unnec- 
essary and  disadvantageous.  Our  experiments  go  to  show  that 
good  results  may  be  expected  when  the  solution  of  the  phosphate 
containing  a  moderate  excess  of  the  magnesium  salt  and  not 
more  than  five  to  ten  per  cent  of  ammonium  chloride  is 
precipitated  by  making  it  slightly  ammoniacal,  the  precipitate 
being  washed  in  slightly  ammoniacal  wash-water.  In  general, 
however,  and  especially  when  more  ammonium  chloride  than 
this  proportion,  or  more  magnesium  salt  than  twice  the  amount 
theoretically  necessary,  is  present,  it  is  safer  to  decant  the 
supernatant  liquid  from  the  precipitate  (through  the  filter  to 

*  Jour.  Am.  Chem.  Soc.,  xx,  445. 


MAGNESIUM  PHOSPHATE   OF  ANALYSIS.          205 

be  used  subsequently  to  hold  the  phosphate),  to  dissolve  the 
precipitate  in  a  little  hydrochloric  acid,  and  reprecipitate  by 
dilute  ammonia,  washing  with  faintly  ammoniacal  wash-water. 
Since  our  first  publication  of  the  work  described  above, 
Neubauer*  has  called  attention  to  the  fact  that  the  influence 
of  ammonia  and  ammonium  salts  upon  the  constitution  of  the 
ammonium  magnesium  phosphate  obtained  in  determining 
phosphoric  acid  had  been  previously  pointed  out  by  him  in  a 
paper  |  discussing  methods  for  the  estimation  of  that  acid. 
We  take  pleasure,  therefore,  in  conceding  to  Neubauer  full 
priority  in  the  observation  of  the  effect  which  we  have 
endeavored  to  overcome  in  the  determination  of  magnesium 
and  of  phosphoric  acid. 

*  Zeitschr.  anorg.  Chem.,  xxii,  162. 
t  Zeitschr.  anorg.  Chem.,  ii,  45. 


XXV 


THE   INFLUENCE    OF    HYDEOCHLOEIC  ACID 
IN    TITEATIONS  BY  SODIUM  THIOSULPHATE  WITH 
SPECIAL  EEFEEENCE  TO  THE  ESTIMA- 
TION OF  SELENIOUS  ACIDS. 

BY  JOHN  T.  NORTON  JR.* 

IN  the  method  of  Norris  and  Fay  f  for  the  iodometric  deter- 
mination of  selenious  acid,  advantage  is  taken  of  a  direct 
and  unique  action  of  sodium  thiosulphate  upon  selenium 
dioxide  in  the  presence  of  hydrochloric  acid.  Most  excellent 
results  are  claimed  for  this  method;  but  the  explicit  state- 
ment of  the  originators  of  the  method,  that  the  amount  of 
hydrochloric  acid  present  does  not  influence  the  result,  pro- 
vided the  titration  is  made  at  the  temperature  of  melting  ice, 
is  so  extraordinary  in  view  of  generally  accepted  ideas  in 
regard  to  the  interaction  of  hydrochloric  acid  and  sodium 
thiosulphate,  as  to, suggest  the  necessity  of  careful  investiga- 
tion of  this  point. 

Pickering,  J  in  his  investigation  of  the  reaction  between 
iodine  and  sodium  thiosulphate,  has  shown  that  more  iodine 
is  required  to  oxidize  the  thiosulphate  as  the  proportion  of 
hydrochloric  acid  increases.  He  ascribed  this  effect  to  the 
formation  of  a  sulphate,  apparently,  by  the  increased  activity 
of  the  iodine,  but  the  more  rational  explanation  is  that, 
although  some  sulphate  is  ultimately  formed,  the  thiosul- 
phate is  first  partially  decomposed  into  free  sulphur  and 

*  From  Am.  Jour.,  Sci.  vii,  287. 

t  Am.  Chem.  Jour.,  vol.  xviii,  p.  703. 

t  Jour.  Chem.  Soc.,  vol.  xxxvii,  p.  135. 


HYDROCHLORIC  ACID  IN  TITRATIONS. 


207 


sulphur  dioxide.      Finkener*  and  Mohrf  also  mention  the 
decomposing  effect  of  free  acid  upon  sodium  thiosulphate. 

The  sodium  thiosulphate  used  in  the  following  experiments 
was  taken  in  nearly  decinormal  solution  and  was  standardized 
by  running  it  into  an  approximately  decinormal  solution  of 
iodine,  the  value  of  which  had  been  determined  by  comparison 
with  decinormal  arsenious  acid  made  from  carefully  resub- 
limed  arsenious  oxide.  In  the  experiments  of  Table  I  the 
solutions  were  stirred  continuously  and  kept  at  a  temperature 
of  from  0°  to  5°  C.,  while  the  thiosulphate  ran  into  the 
acidified  liquid.  The  volume  of  the  solution,  though  fixed  at 
the  beginning  as  given  in  the  table  was  considerably  increased 
during  the  operation  by  the  melting  of  the  ice.  Titrations 

TABLE  I. 


Volume  of 
liquid  at 
beginning  of 
titration. 

Na^Og 

approx- 
imately 
^  taken. 

Volume  of  ^  iodine  used  in  titration. 

HC1  =  none. 

=:  1  cms. 

=  5  cm3. 

=  10  cm3. 

cm3 

cm3 

cm» 

cm3 

cm3 

cm8 

100 

30 

30.251 

30.75 

30.76 

31.20 

200 

30 

30.22 

30.21 

30.56 

31.40 

300 

30 

30.20 

mean  = 

on  99 

30.22 

31.03 

30.90 

400 

30 

30.21 

OvwDS 

30.20 

30.20 

30.55 

500 

30 

30.20 

30.20 

30.21 

30.55 

100 

25 

25.29' 

25.32 

25.98 

25.70 

200 

25 

25.28 

25.34 

25.40 

25.45 

300 

25 

25.29 

mean  = 

fc"*C    OT 

25.41 

25.38 

25.83 

400 

25 

25.27 

25.27 

25.24 

25.30 

25.63 

500 

25 

25.22  . 

25.23 

25.40 

25.30 

100 

20 

20.15 

20.17 

20.33 

20.23 

200 

20 

20.20 

20.13 

20.27 

20.23 

300 

20 

20.21 

mean  =: 

9O  1  1\ 

20.15 

20.20 

20.17 

400 

20 

20.20 

Zv.lO 

20.10 

20.27 

20.07 

500 

20 

20.10 

20.10 

20.17 

20.13 

were  conducted  as  rapidly  as  possible  to  avoid  the  separation 
of  sulphur,  which  is  likely  to  occur,  especially  when  the  acid 
and  thiosulphate  are  present  in  large  quantities.  A  perusal 
of  the  table  shows  that  the  influence  of  the  hydrochloric  acid 
upon  the  thiosulphate  depends  chiefly  upon  the  amount  of 


*  Anal.  Chem.,  6.  Aufl.,  p.  620.          t  Titrirmethode,  6.  Aufl.,  p.  279. 


208 


INFLUENCE  OF  HYDROCHLORIC  ACID  IN 


the  thiosulphate  present  and  afterwards  upon  the  degree  of 
dilution  and  its  own  absolute  quantity.  Thus  when  30  cm3 
of  sodium  thiosulphate  were  employed  the  effect  of  10  cm3 
of  acid  is  marked  at  all  dilutions  within  the  range  of  the 
experiments;  the  effect  of  5  cm3  of  acid  is  inappreciable  only 
at  a  dilution  of  from  400  to  500  cm3,  and  when  1  cm3  of  acid  is 
employed  the  effect  is  only  perceptible  at  a  volume  of  100  cm3. 
When  25  cm3  of  the  thiosulphate  were  used  the  influence  of 
the  acid  is  less  marked;  for  at  a  dilution  of  500  cm3  the 
effect  of  10  cm8  of  acid  is  not  seen,  and  20  cm3  of  the  thio- 
sulphate may  be  present  at  any  dilution  down  to  100  cm3  in 
the  presence  of  as  much  as  10  cm3  of  the  acid,  and  even 
considerably  more,  as  experiments  not  included  in  the  table 
indicated. 

The  slight  discrepancies  which  appear  occasionally  in  the 
table  were  due,  no  doubt,  to  unavoidable  differences  in  the 
tune  of  action. 

This  influence  of  time  upon  the  reaction  between  sodium 
thiosulphate,  iodine,  and  hydrochloric  acid  comes  out  clearly  in 
the  following  series  of  experiments,  in  which  the  thiosulphate 
was  run  into  the  acidified  water,  cooled  to  a  temperature  of 
from  0  to  5°  C.  by  means  of  ice,  the  solution  being  allowed 
to  stand  5,  10,  and  15  minutes.  Sulphur  was  thrown  down 
in  nearly  every  case. 

TABLE  IL 


Volume  of 
the  liquid  at 
beginning  of 
titration. 

HC1 
(sp.  gr.  1.12) 
present. 

Na2S,08 
approxi- 
mately. 
£  taken. 

Volume  of  ^  iodine  used  in  titration  after  standing. 

5  minutes. 

10  minutes. 

15  minutes. 

cm8 

cm8 

cm8 

cm8 

cm8 

cm8 

200 

10 

30 

30.80 

31.30 

32.32 

200 

10 

25 

25.50 

26.00 

26.30 

200 

10 

20 

20.30 

20.70 

20.68 

The  results  of  the  table  emphasize  sufficiently  the  necessity 
of  proceeding  as  rapidly  as  possible  with  the  titration  of 
sodium  thiosulphate  by  iodine  in  presence  of  hydrochloric 


TITRATIONS  BY  SODIUM  THIOSULPHATE. 


209 


acid,  when  the  thiosulphate  is  present  in  considerable  amount. 
As  would  be  expected,  the  effect  of  temperature  upon  the 
reaction  is  also  marked.  In  the  following  experiments  the 
sodium  thiosulphate  was  run  into  the  acidified  water,  pre- 
viously heated  to  the  temperature  indicated,  and  then  titrated 

with  iodine. 

TABLE  m. 


Volume  of 

Volume  of  ^ 

liquid  at 
beginning 
of 

HCl 
(sp.  gr.  1.12) 

Temperature 
Centigrade. 

approxi- 
mately. 

iodine  used 
in  titratious  at 
different 

titration. 

is 

temperatures. 

cm3 

cm» 

C. 

cm8 

cm3 

400 

10 

6° 

25 

23.52 

400 

10 

22° 

25 

23.73 

400 

10 

34° 

25 

24.35 

400 

10 

42° 

25 

24.5 

400 

10 

54° 

25 

25 

400 

10 

64° 

25 

26.1 

From  these  results  it  is  plain  that  the  conditions  under  which 
considerable  amounts  of  sodium  thiosulphate  are  titrated  in 
presence  of  hydrochloric  acid  must  be  carefully  guarded 
when  accuracy  is  a  consideration.  It  is  also  apparent  that 
in  all  cases  the  temperature  should  be  reduced  as  nearly  to 
0°  C.  as  possible  and  rapidity  of  titration  by  the  iodine  is 
an  essential.  So  long  as  the  thiosulphate  present  does  not 
exceed  20  cm3  of  the  T^  solution,  rapid  titration  in  cold 
solution  proceeds  with  fair  regularity  in  presence  of  hydro- 
chloric acid  up  to  10  cm3  of  the  acid  of  sp.  gr.  1.12.  When, 
however,  the  amount  of  thiosulphate  is  greater  than  20cm3 
of  the  -fy  solution,  care  as  to  the  restriction  of  the  acid  and 
dilution  of  the  solution  becomes  a  necessity.  Fortunately, 
in  most  analytical  processes  involving  the  use  of  the  thio- 
sulphate it  is  possible  to  add  that  reagent  from  the  burette 
to  the  solution  to  be  acted  upon,  so  that  it  is  destroyed 
normally  as  fast  as  it  is  introduced  and  the  danger  of  inter- 
action with  the  acid  does  not  occur.  In  the  process  of  Norris 
and  Fay,  however,  the  method  involves  the  addition  of  an 
excess  of  the  thiosulphate  to  the  solution  of  selenious  and 

VOL.    II.  —  14 


210         INFLUENCE  OF  HYDROCHLORIC  ACID  IN 

hydrochloric  acids,  and  thus  the  conditions  prevail  which 
demand  care  as  to  the  relation  of  the  acid,  the  thiosulphate 
and  the  degree  of  dilution.  I  have  experimented,  therefore, 
with  this  process  under  varying  conditions. 

The  process  of  Norris  and  Fay  for  the  iodometric  determi- 
nation of  selenious  acid  consists  briefly  in  treating  the  solution 
of  that  acid  in  ice  water,  in  the  presence  of  hydrochloric  acid, 
with  an  excess  of  a  T*ff  solution  of  sodium  thiosulphate  and 
titrating  back  the  excess  of  the  thiosulphate  with  iodine. 
Four  molecules  of  sodium  thiosulphate  act,  apparently,  upon 
one  molecule  of  selenious  acid  according  to  a  reaction  which 
the  authors  propose  to  study. 

The  selenium  dioxide  used  was  made  by  dissolving  presum- 
ably pure  selenium  in  nitric  acid  and  evaporating  to  dryness. 
The  residue  was  then  treated  with  water,  and  a  little  barium 
hydroxide  was  added  to  remove  any  sulphate  which  might  be 
present.  The  solution  was  then  filtered  and  the  filtrate 
evaporated  to  dryness.  The  residue  was  mixed  with  four  or 
five  times  its  volume  of  dried  pulverized  pyrolusite,  and 
the  whole  was  put  into  a  porcelain  crucible  and  heated.  The 
sublimate  of  selenium  dioxide  was  carefully  collected  on  a  dry 
watch-glass  and  put  into  a  drying  bottle  as  quickly  as  possible. 
The  pyrolusite  prevents  any  reduction  of  the  selenium  dioxide 
to  selenium  and  the  product  consisted  of  beautiful  long  white 
needles.  This  method  of  preparing  the  selenium  dioxide, 
which  has  been  used  for  some  time  in  this  laboratory,  avoids 
contamination  of  the  selenium  dioxide  by  nitric  acid  or  water, 
resulting  from  the  decomposition  of  the  latter,  which  would 
be  possible  in  case  this  reagent  were  employed  in  the  final 
sublimation,  as  is  recommended  by  Norris  and  Fay.  The 
hydrochloric  acid  used  was  of  a  sp.  gr.  1.12,  as  recommended 
by  the  originators  of  the  process.  For  the  experiments  of  Table 
IV  the  dilution  at  the  beginning  was  fixed  at  400  cm3,  and  this 
was  increased  in  every  case  by  the  melting  of  the  ice  used  to 
cool  the  liquid.  A  glance  at  the  preceding  part  of  this  paper 
will  show  that  at  this  degree  of  dilution  the  hydrochloric  acid 
present  has  the  least  effect. 


TITRATIONS  BY  SODIUM  THIOSULPHATE.        211 


TABLE  IV. 


Exp. 

Amount 
SeO, 
taken. 

HCl(sp. 
gr.  1.12) 
taken. 

Volume 
at  begin- 
ning of 
titration. 

Excess 
Na,S203 
employed. 

SeO,  found. 

Error. 

grin. 

cm3 

cm8 

cm3 

grm. 

grm. 

(1) 

0.0616 

10 

400 

2.28 

0.0625 

0.0009+ 

(2) 

0.0628 

10 

400 

7.11 

0.0631 

0.0003+ 

(3) 

0.0508 

10 

400 

11.4 

0.0511 

0.0003+ 

(4) 

0.0587 

10 

400 

12.8 

0.0594 

0.0007+    mean. 

(5) 
(6) 

0.0807 
0.0633 

10 
10 

400 
400 

15.3 

20.85 

0.0813 
0.0638 

0.0006+  0.00005+ 
0.0005+ 

(7) 

0.0682 

25 

400 

1.11 

0.0685 

0.0003+ 

(8) 

0.0779 

25 

400 

1.35 

0.0788 

0.0009+ 

(9) 

0.0465 

25 

400 

18.93 

0.0469 

0.0004+ 

These  results,  while  not  so  good  as  those  obtained  by  Norris 
and  Fay,  are  satisfactory  and  show  that  at  this  degree  of 
dilution  the  process  is  accurate.  These  results  accord  closely 
with  those  contained  in  Table  I.  At  a  dilution  of  400  cm3  or 
in  the  presence  of  only  20  cm3  of  sodium  thiosulphate  in 
excess  the  hydrochloric  acid  present  had  no  perceptible  effect. 
Of  course,  it  must  be  kept  in  mind  that  the  hydrochloric  acid 
acts  only  upon  the  excess  of  thiosulphate  which  is  not  taken 
up  by  the  selenium  dioxide.  The  slight  constant  plus  error 
in  these  results  cannot  be  accounted  for  by  errors  in  the 
standards;  they  were  all  carefully  determined.  Another 
preparation  of  selenium  dioxide  was  made,  starting  with  pure 
selenium  carefully  precipitated  by  sulphurous  acid,  before 
putting  it  through  the  course  of  treatment  previously  described, 
and  the  results  obtained  by  the  action  of  the  sodium 
thiosulphate  recorded  in  Table  V  agree  closely  with  those 
of  the  preceding  table. 

TABLE  V. 


Exp. 

Amount 
SeO, 
taken. 

HC1 
(sp.  gr. 
1.12.) 

H?Oat 
beginning. 

NajSjOg  in 
excess. 

SeO,  found. 

Error. 

grm. 

cm* 

cm8 

cm8 

grm. 

grm. 

(1) 

0.0562 

10 

400 

9.52 

0.0566 

0.0004+ 

(2) 

0.0651 

25 

400 

11.20 

0.0655 

0.0004+ 

The  next  step  was  to  determine  the  effect  of  diminishing  the 


212        INFLUENCE  OF  HYDROCHLORIC  ACID  IN 


dilution  and  of  varying  the  strength  of  acid.    The  following 
table  gives  the  results  of  my  experiments. 


TABLE  VI. 


Exp. 

Amount 
of  SeO2 
taken. 

Volume 
of  H2O 
at  begin- 
ning. 

HC1 

(sp.  gr. 
i:i2). 

Excess  of 
NajSjOj. 

SeO, 
taken. 

Error. 

grin. 

cm3 

cm3 

cm* 

grm. 

grm. 

(1) 

0.1042 

200 

5 

24.16 

0.1041 

0.0001- 

(2) 

0.0611 

200 

10 

13.3 

0.0611 

0.0000 

(3) 

0.0850 

200 

10 

21.9 

0.0828 

0.0022- 

(4) 

0.0757 

200 

25 

13.07 

0.0749 

0.0008- 

(5) 

0.0540 

200 

25 

21.02 

0.0522 

0.0018- 

(6) 

0.0674 

300 

5 

10.04 

0.0679 

0.0005-1- 

(7) 

0.2416 

400 

5 

15.9 

0.2424 

0.0008+ 

It  is  apparent  that  at  the  dilution  of  200  cm3  we  run  into 
difficulties,  and  the  greater  the  excess  of  thiosulphate  present 
the  greater  is  the  error.  When  the  amount  of  sodium 
thiosulphate  exceeds  20  cm3  a  reduction  in  the  amount  of  acid 
to  5  cm3  is  plainly  of  advantage,  as  is  shown  in  a  comparison 
of  Exps.  (1),  (3),  and  (5),  and  is  not  disadvantageous  at 
larger  dilutions  and  with  smaller  amounts  of  the  thiosulphate, 
as  shown  in  Exps.  (6)  and  (7).  The  necessity  of  placing 
some  limits  on  the  method  of  Norris  and  Fay  has  now,  I 
think,  been  established.  The  excess  of  the  thiosulphate  must 
be  carefully  regulated,  as  well  as  the  temperature.  If  one  has 
knowledge  of  the  approximate  amount  of  selenious  acid  in 
solution,  this  is  not  a  matter  of  great  difficulty,  and  things 
should  be  so  arranged  that  no  more  than  20  cm3  of  the  ^ 
thiosulphate  should  ever  be  present  in  excess  of  the  amount 
necessary  to  reduce  the  selenious  acid.  If  this  limit  — 
amounting  to  0.0400  cm8  of  SeO2  —  is  placed  upon  the 
thiosulphate,  so  much  as  10  cm3  of  hydrochloric  acid  (sp.  gr. 
1.12)  may  be  present  without  endangering  the  accuracy  of  the 
process,  provided  the  solution  is  diluted  to  400  cm3  at  the 
outset;  if  only  5  cm8  of  hydrochloric  acid  are  present, 
the  volume  at  the  beginning  may  be  reduced  with  safety  to 
200  cm3.  At  all  events,  5  cm3  of  the  hydrochloric  acid  are 


TITRATIONS  BY  SODIUM  THIOSULPHATE.        213 


amply  sufficient  to  bring  about  the  reaction  between  the 
thiosulphate  and  the  selenium  at  any  dilution  within  the  range 
of  my  experiments.  With  these  precautions  taken,  the  process 
of  Norris  and  Fay  is  simple,  rapid,  and  accurate ;  without  them, 
as  the  experimental  results  indicate,  errors  of  considerable 
amount  may  enter. 

According  to  the  method  of  Muthmann  and  Shafer,*  the 
determination  of  selenious  acid  is  effected  by  the  simple 
addition  of  potassium  iodide  to  the  acidulated  solution  of 
selenious  acid,  and  the  iodine  set  free  is  titrated  with  sodium 
thiosulphate.  In  this  procedure  the  thiosulphate  is  taken  up 
by  the  iodine  as  it  is  added  to  the  solution,  so  that  the  danger 
of  any  action  between  the  thiosulphate  and  the  acid  is  out  of 
the  question.  It  was  shown  in  a  former  paper  from  this 
laboratory  f  that  this  simple  method  is  inaccurate  on  account 
of  the  incompleteness  of  reduction  in  the  cold  and  in  presence 
of  the  iodine  evolved.  In  a  later  article  also  from  this 
laboratory  J  it  was  shown  that  selenium  may  be  completely 
precipitated  and  determined  with  accuracy  gravimetrically 
provided  the  amount  of  potassium  iodide  employed  is  enor- 
mously in  excess  of  that  theoretically  required.  This  suggests 
naturally  the  trial  of  very  large  excesses  of  potassium  iodide 
in  the  procees  of  Muthmann  and  Shafer.  The  details  of 
experiments  made  in  this  manner  are  given  in  the  following 
table. 

TABLE  VII 


Vol.  of 

HCl 

Exp. 

sa 

KI. 

solu- 
tion. 

(sp.  gr. 
1.12). 

found. 

Error. 

grm. 

grlii. 

cm3 

cm3 

grui. 

grm. 

(1) 

0.0553 

10 

150 

10 

0.0558 

0.0005+ 

(2) 

0.0574 

5 

150 

10 

0.0567 

0.0007- 

(3) 

0.0683 

5 

150 

10 

0.0683 

0.0000 

(4) 

0.0487 

5 

150 

10 

0.0484 

0.0003- 

(5) 

0.2617 

10 

150 

10 

0.2589 

0.0028- 

*  Ber.  Dtsch.  chem.  Ges.,  xxyi,  1008. 

t  Gooch  and  Reynolds,  Am.  Jour.  Sci.,  1,  254.    Volume  I,  p.  310. 

t  Peirce,  Am.  Jour.  Sci.,  i,  1896,  416.    Volume  I,  p.  365. 


214  HYDROCHLORIC  ACID  IN  TITRATIONS. 

It  is  obvious  that  for  small  quantities  of  selenium  dioxide  the 
accuracy  of  the  process  is  very  much  increased  by  the  use  of 
large  amounts  of  iodide,  though,  of  course,  the  difficulty  in 
reading  the  end  reaction  due  to  the  presence  of  precipitated 
red  selenium  still  remains ;  but  the  process  is  still  inaccurate 
when  large  amounts  of  selenium  dioxide  are  employed. 


XXVI 

THE  VOLATILIZATION  OF  THE  IRON  CHLOR- 
IDES IN  ANALYSIS  AND  THE  SEPARATION 
OF  THE  OXIDES  OF  IRON  AND  ALUMINUM. 

BY  F.  A.  GOOCH  AND  FRANKE  STUART  HAVENS  * 

IT  is  well  known  that  metallic  iron  is  easily  acted  upon  by  an 
excess  of  chlorine  at  moderately  elevated  temperatures  with 
the  formation  of  ferric  chloride,  and  that  the  product  of  the 
action  of  hydrochloric  acid  gas  upon  the  metal  is  ferrous 
chloride.  Out  of  contact  with  air,  or  moisture,  both  chlorides 
may  be  volatilized  at  appropriate  temperatures  —  the  ferric 
chloride  below  200°  C. ;  the  ferrous  chloride  at  a  bright  red 
heat.  If  water  vapor,  or  oxygen,  or  air  be  present  during 
the  heating,  both  chlorides  are  partially  decomposed  with 
the  formation  of  non-volatile  residues,  ferric  oxide  or  ferric 
oxy-chloride. 

Analytical  processes  involving  the  volatilization  of  iron  at 
temperatures  more  or  less  elevated,  in  an  atmosphere  of  chlo- 
rine or  hydrochloric  acid,  have  been  the  object  of  considerable  at- 
tention. Thus,  Fresenius,f  Drown  and  Shimer,J  and  Watts,§ 
have  heated  crude  iron  in  chlorine  to  remove  the  metal  and 
leave  the  non-volatile  constitutents ;  and  Sainte-Claire  Deville  || 
has  employed  hydrochloric  acid  to  volatilize  iron  from  mix- 
tures of  that  metal  with  alumina  (obtained  by  heating  the 
mixed  oxides  of  iron  and  aluminum  in  hydrogen  according  to 
Rivot),^[  exposing  the  mixture,  contained  in  a  porcelain  boat 
and  placed  within  a  porcelain  tube,  to  the  bright  red  heat 
of  a  charcoal  furnace  —  an  operation  which  was  bettered  by 

*  From  Am.  Jour.  Sci.,  yii,  370.  t  Zeitschr.  anal.  Chem.,  iv,  72. 

t  Jour.  Inst.  Min.  Eng.,  viii,  613.  §  Chem.  News,  xlv,  279. 

||  Ann.  China.  [3],  xxxviii,  23.  IT  Ann.  Chim.  [3],  xxx,  188. 


216  THE   VOLATILIZATION  OF  THE 

Cooke's*  use  of  a  tube  of  platinum  instead  of  the  porcelain 
tube  and  a  gas  blowpipe  in  place  of  the  charcoal  furnace. 
Sainte-Claire  Devillef  showed,  further,  that  ferric  oxide  may 
be  converted  to  ferric  chloride  and  volatilized  at  the  heat  of 
the  charcoal  furnace  if  the  current  of  hydrochloric  acid  is 
sufficiently  rapid ;  but  the  curious  effect  was  observed  that  in 
a  sufficiently  limited  current  of  the  acid  no  chloride  whatever 
was  volatilized,  while  the  amorphous  oxide  was  converted  to 
the  highly  crystalline  oxide  of  the  same  composition — a  phe- 
nomenon which  gave  rise  to  a  theory  of  the  natural  formation 
of  specular  iron  in  volcanic  regions. 

Quite  recently,  Mover  £  has  made  record  of  an  unsuc- 
cessful attempt  (in  the  course  of  experimentation  upon  the 
volatility  of  certain  chlorides  at  comparatively  low  tempera- 
tures) to  convert  ferric  oxide  completely  to  ferric  chloride  by 
the  action  of  gaseous  hydrochloric  acid  at  about  200°  C.  At 
this  temperature  the  greater  part  of  the  iron  sublimed,  but  a 
residue  remained,  which,  volatilizing  neither  on  long  heating 
at  200°  nor  upon  considerable  elevation  of  the  temperature, 
proved  upon  examination  to  be  ferrous  chloride.  In  the 
experiments  to  be  described  we  have  acted  with  gaseous 
hydrochloric  acid  upon  ferric  oxide  made  by  igniting  the 
nitrate  prepared  from  pure  iron  deposited  electrolytically  by 
high  currents  passing  between  electrodes  of  platinum  in  a 
strong  solution  of  ammonio-ferrous  sulphate.  The  oxide, 
contained  in  a  porcelain  boat,  was  heated  within  a  roomy  glass 
tube  over  a  small  combustion  furnace.  The  hydrochloric  acid 
(generated  by  dropping  sulphuric  acid  into  a  mixture  of 
strong  hydrochloric  acid  and  salt,  and  dried  by  calcium  chlor- 
ide) entered  one  end  of  the  tube  and  passed  out  at  the  other 
through  a  water  trap.  In  early  experiments  a  high-reading 
thermometer  was  inserted  through  the  stopper  in  the  exit 
end  of  the  tube  so  that  its  bulb  was  above  and  immediately 
adjacent  to  the  boat  carrying  the  oxide.  In  this  way  the 
actual  temperatures  of  the  vapors  about  the  boat  were  fixed 

*  Am.  Jour.  Sci.,  xlii,  78.  t  Compt.  rend.,  Hi,  1264. 

J  Jour.  Am.  Chem.  Soc.,  xviii,  1029. 


IRON  CHLORIDES  IN  ANALYSIS.  217 

with  considerable  accuracy ;  later,  after  a  little  experience  in 
gauging  the  effect  of  the  burners,  it  was  found  that  the  tem- 
peratures could  be  regulated  very  closely  without  actually 
depending  upon  the  thermometer.  We  found,  as  did  Moyer, 
that  ferric  oxide,  submitted  to  the  action  of  dry  hydrochloric 
acid  gas,  volatilizes  partially  as  ferric  chloride  at  low  tempera- 
tures—180°  to  200°  C.  — leaving  ultimately  a  crystalline 
residue  which  does  not  change  visibly  when  heated  for  an 
hour  or  two  at  200°,  or  even  at  500°,  in  the  pure  dry  acid. 
According  to  our  experience,  this  residue  is  generally  slightly 
reddish  or  salmon-colored ;  but  sometimes,  especially  after  a 
second  heating,  the  boat  having  been  withdrawn  from  the 
tube  or  exposed  to  the  atmosphere  (and  so  to  moisture),  the 
residue  is  white.  When  it  is  white  it  dissolves  in  water, 
yields  'the  characteristic  reaction  for  a  ferrous  salt  with  potas- 
sium ferricyanide,  gives  no  reaction  with  potassium  sulphocy- 
anide,  and  upon  treatment  in  weighed  amount  with  potassium 
permanganate  destroys  the  amount  of  that  reagent  theoretically 
required  for  its  oxidation  upon  the  supposition  that  it  is  fer- 
rous chloride.  The  slightly  colored  residue  when  treated  with 
water  yields  a  solution  showing  the  reaction  of  a  ferrous  salt 
only,  but  when  treated  with  hydrochloric  acid  and  then  tested 
shows  the  presence  of  a  trace  of  iron  hi  the  ferric  condition. 
Doubtless  the  coloration  of  the  residue  is  due  to  an  included 
trace  of  ferric  oxide  or  oxychloride,  which  after  exposure  of 
the  containing  crystals  to  slight  atmospheric  action,  is  more 
easily  reached  in  the  second  heating  by  the  gaseous  acid.  The 
amount  of  residue  is  somewhat  variable,  but  approximates 
under  the  conditions  of  our  work  to  from  five  to  ten  per  cent 
of  the  oxide  taken :  thus,  in  one  typical  experiment  0.1  grm. 
of  ferric  oxide  left  a  residue  which  (withdrawn  after  cooling) 
weighed  0.0115  grm. 

The  greater  portion  of  the  ferric  oxide  volatilizes  when 
submitted  to  the  action  of  the  gaseous  acid  at  200°  quickly 
and  abundantly  in  the  form  of  the  greenish  vapor  of  ferric 
chloride,  and  if  the  operation  is  interrupted  at  this  stage  the 
residue  which  remains  is  nearly  black,  insoluble  in  water, 


218  THE   VOLATILIZATION  OF  THE 

slightly  soluble  in  cold  hydrochloric  acid,  and  readily  soluble 
in  hot  hydrochloric  acid  with  the  formation  of  ferric  chloride. 
It  is  probably  something  analogous  to  the  oxychloride  which 
Rousseau*  identifies  as  the  product  of  the  action  of  water 
upon  ferric  chloride  at  275°  to  300°.  This  dark  residue 
yields  to  the  action  of  the  hydrochloric  acid  at  180°  to  200° 
only  slowly,  but  ultimately  only  the  residue  which  is  essen- 
tially ferrous  chloride  remains ;  thereafter  little  volatilization 
occurs  within  the  range  of  temperature  of  our  experimentation 
—  200°  to  500°. 

It  is  obvious  that  a  reduction  of  iron  in  the  ferric  condition 
to  iron  in  the  ferrous  condition  takes  place  under  the  con- 
ditions of  our  work,  and  it  in  difficult  to  see  how  this  can 
occur  otherwise  than  by  the  direct  dissociation  of  ferric 
chloride  under  the  low  partial  pressure  conditioned  by  the 
brisk  current  of  hydrochloric  acid  gas.  The  temperature  of 
formation,  180°  to  200°,  is  far  below  that  at  which  such 
dissociation  is  supposed  to  begin.  Thus,  Gruenewald  and 
Meyer  f  found,  after  cooling,  no  evidence  of  the  dissociation 
of  ferric  chloride  which  had  been  heated  in  the  Victor  Meyer 
vapor-density  apparatus  to  448°  in  contact  or  partial  mixture 
with  nitrogen;  but  ten  per  cent  of  the  residue  obtained  by 
heating  to  518°  was  in  the  ferrous  condition.  Friedel  and 
Crafts,!  however,  did  see  crystals  of  ferrous  chloride  at  440° 
on  the  walls  of  a  Dumas  container  filled  with  the  vapor  of 
ferric  chloride  and  nitrogen,  the  former  exerting  a  partial 
pressure  of  0.75;  while  ferric  chloride  volatilized  into  an 
atmosphere  of  chlorine  without  evidence  of  dissociation.  It 
seems  rather  surprising,  therefore,  to  find  so  large  a  percen- 
tage of  dissociation  as  that  shown  in  our  experiments  at  a 
temperature  so  low  — 180°  to  200°.  Curiously,  too,  we  find, 
on  repeating  the  experiment  of  heating  ferric  oxide  in  gaseous 
hydrochloric  acid,  that  if  the  temperature  of  the  oxide  is 
450°  to  500°  when  the  brisk  current  of  acid  begins  to  act, 
the  whole  mass  of  oxide  is  converted  and  volatilizes  without 

*  Compt.  rend.,  cxvi,  118.  t  Ber.  Dtsch.  chem.  Ges.,  xxi,  687. 

|  Compt.  rend.,  cvii,  301. 


IRON  CHLORIDES  IN  ANALYSIS.  219 

residue.  It  is  hardly  to  be  supposed  that  the  degree  of 
dissociation  at  450°  to  500°  can  be  less  than  that  at  180°  to 
200°,  and  a  test  of  the  sublimate,  after  cooling,  shows  that 
it  contains  a  ferrous  salt.  Plainly,  ferrous  chloride  (formed 
by  dissociation)  has  volatilized,  and  inasmuch  as  the  ferrous 
chloride  constituting  the  residue  formed  at  180°  to  200°  does 
not  volatilize  in  the  hydrochloric  acid  even  at  500°  it  is  plain 
that  the  volatility  of  the  former  is  not  determined  by  the 
presence  of  the  latter.  Apparently,  the  cause  of  the  com- 
pleteness of  volatilization  must  be  sought  in  its  rapidity ;  and 
this  is  not  an  unreasonable  hypothesis,  if  one  considers  that 
an  action  sufficiently  rapid  to  keep  above  the  boat  an  atmos- 
phere of  ferric  chloride  and  its  products  of  partial  dissociation, 
might  naturally  provide  the  very  condition  which  would  be 
effective  in  counteracting  the  tendency  of  the  residue  to 
dissociate  before  it  volatilizes.  If  this  hypothesis  is  correct, 
it  is  plain  that  the  introduction  of  chlorine  gas,  the  active 
product  of  dissociation,  into  the  atmosphere  of  hydrochloric 
acid  ought  to  bring  about  the  volatilization  of  the  residue 
of  ferrous  chloride,  formed  at  180°  to  200°,  which  refuses 
to  volatilize  in  the  acid  alone.  As  a  matter  of  fact,  we  find 
by  experiment  that  if  a  little  manganese  dioxide  is  added 
to  the  contents  of  the  generator,  so  that  the  hydrochloric 
acid  may  carry  with  it  a  little  chlorine,  every  trace  of  ferric 
oxide  is  volatilized  from  the  boat  at  180°  to  200°;  and 
the  residue  of  ferrous  chloride  found  at  180°  to  200°  when 
the  hydrochloric  acid  is  used  alone  is  likewise  volatilized 
at  the  same  temperature,  when  the  admixture  of  chlorine 
is  made. 

These  facts,  that  ferric  oxide  is  completely  volatile  hi 
hydrochloric  acid  gas  applied  at  once  at  a  temperature  of 
450°  to  500°  C.,  and  at  180°  to  200°  if  the  acid  carries  a 
little  chlorine,  open  the  way  to  many  analytical  separations 
of  iron  from  substances  not  volatile  under  these  conditions. 
In  the  experiments  of  the  following  table  we  have  applied 
these  methods  to  the  separation  of  intermixed  iron  and 
aluminum  oxides.  The  ferric  oxide  employed  was  made,  as 


220 


THE   VOLATILIZATION  OF  THE 


before,  by  ignition  of  the  nitrate  prepared  from  iron  deposited 
electrolytically  by  a  strong  current  passing  between  platinum 
electrodes  in  a  solution  of  ammonio-ferrous  sulphate.*  The 
aluminum  oxide  was  made  by  igniting  to  a  constant  weight 
the  carefully  washed  hydroxide  precipitated  by  ammonia  from 
a  pure  hydrous  chloride  thrown  down  from  the  solution  of 
a  commercially  pure  chloride  by  hydrochloric  acid.f  The 
hydrochloric  acid  gas  was  made  by  dropping  sulphuric  acid 
into  strong  hydrochloric  acid  mixed  with  salt,  and  a  little 
manganese  dioxide  was  added  when  the  mixture  with  chlorine 
was  desired.  The  experimental  details  are  given  in  the 
table. 


Feo03 
taken. 

A1203 
taken. 

found. 

Error. 

Time. 

Temperature. 

Atmosphere. 

grm. 
0.1000 
0.2000 
0.1020 
0.2145 

gnu. 

o!ibi5 

0.1006 

grm. 

o!ibi5 

0.1008 

grm* 
0.0000 
0.0000 
0.0000 
0.0002+ 

hrs. 

1 

C°. 
450-600 
450-500 
450-500 
450-500 

HC1 
HC1 
HC1 
HC1 

0.1000 
0.1000 
0.1072 
0.2045 
0.1050 
0.2008 

o!l032 
0.1013 
0.1032 
0.1023 
0.1007 
0.1087 

o!l032 

0.1015 
0.1033 
0.1019 
0.1006 
0.1087 

0.0000 
0.0000 
0.0002+ 
0.0001+ 
0.0004- 
0.0001- 
0.0000 

1 

H 
1 

180-200 
180-200 
180-200 
180-200 
450-500 
450-500 
450-500 

HC1  +  CL. 
HC1  +  C12. 
HC1  +  C12. 
HCi  +  C12. 
HC1  +  CL. 
HCI  +  C12. 
HCI  +  C12. 

The  residual  alumina  tested  in  several  experiments  by 
fusion  with  sodium  carbonate,  solution  in  hydrochloric  acid, 
and  addition  of  potassium  sulphocyanide  gave  no  indication 
of  the  presence  of  iron. 

The  separation  of  the  iron  is  obviously  complete  at  450° 
to  500°  when  the  mixed  oxides  are  submitted  at  once  to 
the  action  of  hydrochloric  acid  gas,  or  at  180°  to  200°  when 

*  The  use  of  an  anode  of  commercially  pure  iron  wire  naturally  facilitates 
the  operation,  but  in  our  experience  the  deposit  thus  obtained  is  likely  to 
carry  traces  of  impurity.  In  an  attempt,  too,  to  prepare  pure  ferric  oxide 
from  the  oxalate  thrown  down  out  of  ferrous  sulphate  with  all  precautions, 
the  material  obtained  still  held  traces  of  silica,  and  possibly  alumina,  amount- 
ing to  0.0004  grm.  in  0.1  grm.  of  the  oxide. 

t  From  Am.  Jour.  Sci.,  ii,  416.    This  volume,  p.  20. 


IRON  CHLORIDES  IN  ANALYSIS.  221 

chlorine  is  mixed  with  the  hydrochloric  acid.  Plainly,  the 
extremely  high  temperatures  employed  by  Deville  are  un- 
necessary if  the  mixed  oxides  are  submitted  at  once  to  the 
action  of  hydrochloric  acid  at  450°  to  500°  without  previous 
gentle  heating  in  the  acid  atmosphere.  We  prefer,  however, 
to  use  the  mixture  of  chlorine  and  hydrochloric  acid,  not 
only  because  the  temperature  of  the  reaction  is  lower,  but 
because  it  needs  no  regulation,  while  the  danger  of  error 
arising  from  the  liability  of  ferric  chloride  to  dissociate,  or 
from  deficiency  of  oxidation  in  the  oxide  treated,  or  from 
mechanical  loss  due  to  rapid  volatilization,  is  avoided. 


XXVII 

THE  TITRATION   OF    OXALIC  ACID    BY  POTAS- 
SIUM PERMANGANATE   IN  PRESENCE   OF 
HYDROCHLORIC  ACID. 

BY  F.  A.  GOOCH  AND  C.  A.  PETEES.* 

LOWENTHAL  and  LENSSEN  f  were  the  first  to  show  that  the 
titration  of  a  ferrous  salt  by  potassium  permanganate  in  the 
presence  of  hydrochloric  acid,  according  to  the  process  of 
Margueritte  J  is  vitiated  by  the  evolution  of  chloride  outside 
the  main  reaction,  and  to  point  out  that  a  remedy  for  the 
difficulty  is  to  be  found  in  the  titration  of  the  ferrous 
salt  in  divided  portions,  other  equal  volumes  of  the  ferrous 
solution  being  added  to  the  liquid  in  which  the  first  titration  is 
accomplished  until  the  amount  of  iron  indicated  by  successive 
titrations  becomes  constant. 

Kessler§  showed  the  restraining  influence  of  certain  sulphates, 
of  manganous  sulphate  in  particular,  upon  the  irregular  and 
undesirable  interaction  of  the  permanganate  and  hydrochloric 
acid,  and  Zimrnermann,  ||  in  apparent  ignorance  of  Kessler's 
forgotten  proposal,  advocated  the  introduction  of  a  manganous 
salt,  best  the  sulphate,  into  the  ferrous  salt  to  be  determined, 
thus  accomplishing  the  purpose  of  the  empirical  procedure  of 
Lowenthal  and  Lenssen. 

The  tendency  toward  evolution  of  chlorine  in  the  oxidation 
of  a  ferrous  salt  by  permanganate,  as  compared  with  the 
absence  of  such  tendency  in  the  similar  oxidation  of  oxalic 
acid,  in  presence  of  hydrochloric  acid,  was  explained  by 

*  From  Am.  Jour.  Sci.,  vii,  461.  t  Zeitschr.  anal.  Chem.,  i,  329. 

t  Ann.  Chim.  Phys.  [3],  xviii,  244. 

§  Ann.  Phys.,  cxciv,  48  (1863) ;  CXCT,  225  (1863). 

|i  Ann.  Chem.,  ccxiii,  302. 


TITRATION  OF  OXALIC  ACID,  ETC.  223 

Zimmermann  on  the  hypothesis  that  an  oxide  of  iron  higher 
than  ferric  oxide  is  formed  as  an  intermediate  product,  and 
that  this  unstable  oxide  is  sufficiently  active  to  break  up 
hydrochloric  acid  as  well  as  to  oxidize  more  of  the  ferrous 
salt.  Quite  recently,  Wagner  *  finds  explanation  of  the 
sensitiveness  of  the  hydrochloric  acid  solution  of  the  ferrous 
salt  in  the  probable  formation  of  chlor-f errous  acid  (analogous 
to  chlor-platinic  and  chlor-auric  acids),  which  suffers  oxidation 
more  readily  than  hydrochloric  acid  under  the  action  of  the 
permanganate.  The  protective  influence  of  the  manganous 
salt  turns  apparently,  as  Zimmermann  suggested,  upon  the 
initiation  of  Guyard's  reaction,  according  to  which  the  per- 
manganate and  manganous  salt  interact  to  form  a  higher 
oxide  of  manganese  of  a  constitution  approaching  the  dioxide 
more  or  less  closely  —  this  oxide  being  capable  of  oxidizing 
the  ferrous  salt,  but  slow  to  act  upon  the  hydrochloric  acid, 
or  the  chlor-ferrous  acid  of  Wagner.  According  to  Volhard,f 
the  reaction  of  Guyard  is  favored  and  hastened  by  heat  and 
concentration  of  the  solution,  while  it  is  delayed  by  acidity 
and  dilution ;  but  even  in  solutions  containing  very  little 
manganous  salt  and  a  considerable  quantity  of  free  acid  the 
faint  rose  color  developed  by  the  careful  addition  of  perman- 
ganate ultimately  vanishes  until  every  trace  of  the  manganous 
salt  is  precipitated.  When  a  considerable  amount  of  the  salt 
is  present  interaction  follows  immediately  the  introduction  of 
the  permanganate.  Zimmermann  advocates  the  use  of  4  grams 
of  manganous  sulphate  uniformly  in  titrations  of  a  ferrous 
salt  by  permanganate,  a  procedure  to  which  Wagner  gives 
acquiescence,  though  pointing  out  that  a  ninth  of  that  amount 
is  all  that  he  finds  to  be  necessary.  The  excess  of  the 
manganous  salt  can  do  no  harm  so  long  as  the  higher  oxide, 
the  product  of  interaction  of  the  manganous  salt  and  the 
permanganate,  is  immediately  reduced  by  even  traces  of  a 
ferrous  salt,  and  this  appears  to  be  the  case  at  least  within 
the  limits  proposed  by  Zimmermann  and  Wagner.  Thus  we 

*  Massanalytische  Studien,  Habilitationsschrift,  Leipzig,  1898. 
t  Ann.  Chem.,  cxcviii,  318,  1879. 


224 


TITRATION  OF  OXALIC  ACID 


find,  as  shown  in  results  of  the  accompanying  table,  that  so 
much  as  five  grams  of  the  sulphate  may  be  present  in  135  cm3 
of  the  liquid,  containing  about  5  cm3  of  hydrochloric  acid  of 
full  strength,  without  interfering  with  the  regularity  of  the 
titration ;  and  the  effect  is  trivial  even  when  the  amount  of 
manganous  sulphate  reaches  ten  grams.  We  find  also  practical 
regularity  of  working  when  manganous  chloride  is  substituted 
for  the  sulphate,  and  in  this  respect  our  results  accord  with 
those  of  Zimmermann  and  differ  from  those  of  Wagner.* 


Total  volume 
at  beginning 
of  titration. 

HCl 

(sp.gr.  1.09). 

FeCl,. 

KMn04^. 

MnSO4  .  5H2O. 

MnCl2.4H,O. 

cm3 

cm8 

cm8 

cm8 

grams. 

grams. 

135 

10 

25 

21.70 

1 

135 

10 

25 

21.70 

3 

( 

135 

10 

25 

21.70 

5 

135 

10 

25 

21.75 

7 

§ 

135 

10 

25 

21.75 

10 

. 

145 

20 

25 

21.75 

,     10 

p 

175 

50 

25 

21.75 

10 

135 

10 

25 

21.70 

t 

1 

135 

10 

25 

21.70 

m 

2 

145 

20 

25 

21.70 

t 

2 

155 

30 

25 

21.76 

3 

165 

40 

25 

21.70 

• 

4 

In  all  cases,  however,  in  which  the  larger  amounts  of  manga- 
nous salt  are  present,  the  end  reaction  is  marked  by  the  advent 
of  a  brownish-red  precipitate  rather  than  the  clear  pink  of 
the  soluble  permanganate,  and  it  is  obvious  that  in  case  the 
solutions  to  be  oxidized  were  not  active  enough  to  act  with 
rapidity  upon  the  product  of  the  Guyard  reaction,  difficulty 
might  follow  the  failure  to  adjust  the  conditions  more 
particularly. 

It  has  been  stated  by  Fleischerf  and  Zimmermann  J  that 
hydrochloric  acid  interferes  in  no  way  with  the  titration  of 
oxalic  acid  by  potassium  permanganate.  This  statement, 
however,  is  not  in  accord  with  our  experience;  for  we  find 
that  in  such  titrations  there  is  a  small  though  real  waste  of 

*  Loc.  cit,  p.  104.  t  Volumetric  Analysis ;  trans,  by  Muir,  p.  71. 

J  Loc.  cit. 


BY  POTASSIUM  PERMANGANATE. 


225 


permanganate  proportionate  to  the  amount  of  hydrochloric 
acid  present.  This  fact  is  brought  out  clearly  in  the  comparison 
of  experiments  of  section  A  in  the  following  table,  in  which 
no  hydrochloric  acid  was  present,  with  experiments  B  in  which 
hydrochloric  acid  was  present. 

Temperature  at  beginning,  about  80°  C. 


Approximate 
volume  at 
beginning  of 

titration. 

H,804  1  :  1. 

HCl 
(sp.  gr.  1.09.) 

Ammonium 

oxalate 

* 

KMn04. 

Variation  from 
mean  of  A 
taken  as 
standard. 

A.. 

200 
200 
200 
200 
200 
200 

cm8 
5 
5 
10 
10 
25 
25 

cm» 

cm3 
50 
50 
50 
50 
50 
50 

cm« 
47.50 
47.50 
47.50 
47.50 
47.50 
47.50 

cm3 
0.00 
0.00 
0.00 
0.00 
0.00 
0.00 

] 

B. 

f  150 
I  150 
1  150 
1  150 
(500 
}500 
(500 

10 
10 
10 
10 
5 
10 
10 

2.5 
2.5 
5.0 
10.0 

ib.b 

10.0 

25 
25 
25 
25 
25 
25 
25 

23.80 
23.90 
23.90 
24.00 
23.80 
24.00 
24.10 

0.05+ 
0.15+ 
0.15+ 
0.25+ 
0.05+ 
0.25+ 
0.35+ 

From  these  results  it  is  evident  that,  though  the  error  intro- 
duced by  the  presence  of  the  hydrochloric  acid  during  the 
action  of  the  permanganate  upon  the  oxalic  acid  is  small,  it  is 
plainly  appreciable.  The  questions  arise,  therefore,  first,  as  to 
whether  the  secondary  action  of  the  permanganate  upon  the 
hydrochloric  acid  may  be  prevented  by  the  presence  of  a 
suitable  amount  of  a  manganous  salt,  and,  secondly,  as  to 
whether  in  this  event  the  reducing  agent — the  oxalic  acid  — 
is  sufficiently  active,  like  the  ferrous  salt,  to  prevent  the 
premature  establishment  of  an  end  color  due  to  the  Guyard 
reaction.  The  latter  question  must  naturally  be  settled  before 
the  former  can  be  taken  up.  In  the  accompanying  table  are 
recorded  the  effects  of  varying  amounts  of  manganous  salt  in 

VOL.    IX.  —  15 


226 


TITRATION  OF  OXALIC  ACID 


presence  of  different  amounts  of  sulphuric  acid  in  the  reaction 
of  permanganate  upon  oxalic  acid. 

Temperature  at  beginning,  about  80°  C. 


Volume 
at 
beginning. 

HjSO4  1  :  1. 

Ammonium 
oxalate 

ff 
TV' 

MnS04  .  5H,O. 

KMnO4. 

Variation 
from 

standard. 

cm* 

cm' 

cm* 

grm. 

cm3 

f  130 

5 

25 

§ 

23.75 

0.00 

130 

6 

25 

0.0008 

23.75 

0.00 

130 

5 

25 

0.0032 

23.76 

0.00 

130 

5 

25 

0.0160 

23.75 

0.00 

. 

130 

5 

25 

1 

23.70 

0.05- 

130 

5 

25 

2 

23.75 

0.00 

130 

5 

25 

2.5 

23.60 

0.15- 

130 

5 

25 

3.0 

23.40 

0.25- 

130 

•    5 

25 

4.0 

23.60 

0.15- 

500 

5 

25 

23.80 

0.05+ 

500 

6 

25 

0.0008 

23.80 

0.05-f 

500 

5 

25 

0.0032 

23.80 

0.05+ 

. 

500 

5 

25 

1 

23.70 

0.05- 

500 

5 

25 

2 

23.40 

0.35- 

500 

5 

25 

3 

23.50 

0.25- 

500 

5 

25 

4 

23.30 

0.45- 

130 

no 

25 

1 

23.80 

0.05+ 

130 

ho 

25 

2 

23.75 

0.00 

130 

jio 

25 

3 

23.65 

0.10- 

130 

1  10 

25 

4 

23.50 

0.25- 

130 

(15 

25 

2 

23.75 

0.00 

" 

130 

)l6 

25 

4 

23.70 

0.05- 

130 

1  16 

25 

5 

23.50 

0.25- 

130 

(30 

25 

2 

23.75 

0.00 

130 

?30 

25 

4 

23.70 

0.05- 

130 

<30 

25 

5 

23.75 

0.00 

From  the  results  given  it  is  evident  that  the  persistence  of 
the  Giiyard  reaction  is  liable  to  interfere  with  the  end  reaction 
of  oxidation  of  oxalic  acid  unless  an  adjustment  is  made 
between  the  quantity  of  the  manganous  salt,  the  amount  of 
acid,  and  the  dilution.  In  hot  solutions  of  a  total  volume  of 
130  cm3  at  the  beginning,  no  more  than  2  grms.  of  the 
manganous  sulphate  should  accompany  5  to  10  cm3  of  the 
1 :  1  sulphuric  acid ;  when  the  total  volume  at  the  beginning 
reaches  500  cm3  no  more  than  a  single  gram  of  the  salt  should 
be  present  with  5  cm8  of  the  1  :  1  sulphuric  acid.  The  amount 
of  manganous  salt  may,  however,  be  increased  considerably  if 
the  quantity  of  acid  is  increased. 


OF  THE 

UNIVERSITY 


BY  POTASSIUM  PERMANGANATE. 


227 


As  Kessier  has  noted,  a  sufficiency  of  the  manganous  salt, 
acting  no  doubt  as  the  medium  of  transfer  of  oxygen,  may 
bring  about  interaction  between  the  permanganate  and  the 
oxalic  acid  at  atmospheric  temperatures  without  the  tedious 
delay  ordinarily  encountered  in  the  attempt  to  consummate 
that  action  in  cold  solutions.  It  would  seem  natural  that  the 
manganic  hydroxide  formed  in  the  Guyard  reaction  at  low 
temperatures  should  yield  more  readily  to  the  reducing  action 
of  the  oxalic  acid  than  the  more  anhydrous  form  to  be  expected 
in  hot  solutions,  so  that  at  such  temperatures  the  limits  as  to 
proportions  of  manganous  salt,  acid,  and  dilution,  within 
which  favorable  action  may  take  place,  should  be  wider; 
moreover,  the  undesirable  action  of  the  permanganate  upon 

Temperature  20°-26°  C. 


Number 
of 
experi- 

BMOfc* 

Volume 
at  begin- 
ning of 

titration. 

W 

HCl 

Ammo- 
nium 
oxalate 

KMn04. 

MnSO4  . 
5H20. 

MnCl2. 
4HjO. 

Variation 
from 
standard. 

cm3 

cm' 

cm» 

cm» 

grin. 

grin. 

cm» 

(ij 

130 
130 

10 
10 

25 

25 

23.90 
23.90 

0.0040 
0.0120 

0.15+ 
0.15+ 

(3) 

130 

10 

25 

23.80 

0.0250 

0.05+ 

(4) 

130 

10 

25 

23.75 

0.0400 

0.00 

(5) 

130 

10 

25 

23.76 

0.0500 

0.01+ 

(6) 

130 

10 

25 

23.70 

0.1000 

0.05- 

(7) 

130 

10 

25 

23.75 

0.2000 

0.00 

(8) 

130 

10 

25 

24.20 

0.0200 

0.45+ 

(9) 
(10) 

130 
130 

10 
10 

25 
25 

23.95 
23.80 

0.0200 
0.0400 

0.20+ 
0.05+ 

(11) 

130 

20 

25 

23.75 

0.0400 

0.00 

(12) 

130 

30 

25 

23.75 

0.0400 

0.00 

(13) 

130 

10 

25 

23.75 

1.0000 

t 

0.00 

(14) 

130 

10 

25 

23.75 

2.0000 

> 

0.00 

(15) 

130 

10 

25 

23.75 

3.0000 

0.00 

(16) 

130 

1 

25 

23.72 

1.0000 

0.03- 

(17) 

130 

1 

25 

23.74 

2.0000 

0.01- 

(18) 

130 

1 

25 

23.72 

3.0000 

0.03- 

(19) 

130 

2 

25 

23.70 

0.5000 

0.05- 

(20) 

130 

3 

25 

23.76 

0.5000 

0.00 

Temperature  about  80°. 

(21) 

145 

10 

10 

25 

23.90 

0.5000 

0.15+ 

(22) 

145 

10 

10 

25 

23.70 

1.0000 

0.05- 

(23) 

500 

10 

10 

25 

23.75 

1.0000 

0.00 

(24) 

500 

m 

10 

25 

28.70 

1.0000 

0.05- 

(25) 

500 

•  • 

10 

25 

24.10 

0.5000 

0.35+ 

228  TITRATION  OF  OXALIC  ACID 

hydrochloric  acid,  when  that  acid  is  present,  should  be  less 
appreciable  at  lower  temperatures.  In  our  experiments, 
therefore,  upon  the  oxidation  of  oxalic  acid  by  potassium 
permanganate  in  presence  of  hydrochloric  acid,  we  have  studied 
the  effect  of  varying  the  proportions  of  the  manganous  salt 
both  at  atmospheric  temperatures  and  the  higher  temperatures 
generally  employed. 

From  these  results  it  appears  that  the  presence  of  a  suit- 
able amount  of  manganous  salt — either  the  sulphate  (4-7), 
(13-15),  (22-24)  or  the  chloride  (10-12),  (16-20)  —  is 
capable,  either  in  cold  solution  (1-20)  or  in  hot  solution 
(22-24)  of  preventing  the  action  of  the  permanganate  upon 
the  hydrochloric  acid.  It  appears,  also,  that,  for  a  given 
dilution  and  strength  of  acid,  less  manganous  salt  is  needed 
in  the  cold  solution  (4-7)  than  in  the  hot  solutions  (22-24). 
Thus,  in  the  hot  solution,  at  a  dilution  at  145  cm3  to  500  cm3 
1  grm.  of  manganous  sulphate  must  be  present  with  5  cm3 . 
of  strong  hydrochloric  acid,  with  or  without  sulphuric  acid ; 
while  in  the  cold  solution  0.04  grm.  of  either  the  sulphate 
or  chloride  is  enough  to  secure  adequate  protective  effect. 
Experience  showed,  however,  that  0.5  grm.  or  1.0  grm.  of  the 
manganous  salt  should  be  present  in  order  to  push  the  re- 
action with  reasonable  speed  in  cold  solutions. 

Wagner*  has  made  record  of  the  increased  evolution  of 
chlorine  in  oxidations  of  ferrous  chloride  by  potassium 
permanganate  in  presence  of  various  salts,  of  which  barium 
chloride  was  the  most  active.  We  have  made  some  experi- 
ments, therefore,  to  determine  whether  such  action  would 
appear  in  the  oxidation  of  oxalic  acid  in  cold  solutions 
containing  certain  salts,  and,  if  so,  whether  it  would  be 
preventable  by  the  presence  of  the  manganous  salt  under  our 
conditions  of  working.  From  the  results  given  in  the 
accompanying  table,  it  is  plain  that  the  evolution  of  chlorine 
in  cold  solutions  is  less  in  the  presence  of  these  salts  than 
when  hydrochloric  acid  is  used  without  them,  and  that  such 
evolution  may  be  entirely  prevented  (within  the  proportions 

*  Loc.  cit. 


BY  POTASSIUM  PERMANGANATE. 


229 


of  our  work)  by  the  presence  of  0.5  grm.  to  1  grm.   of 
manganous  chloride. 

Finally,  it  appears  as  the  result  of  an  investigation,  that  the 
titration  of  oxalic  acid  by  potassium  permanganate  in  presence 
of  hydrochloric  acid  is  ordinarily  attended  with  some  inaccu- 
racy due  to  liberation  of  chlorine  from  the  hydrochloric  acid ; 
that  this  tendency  may  be  overcome  by  the  presence  of  a 
manganous  salt — either  the  sulphate  or  chloride;  that  1  grm. 
of  the  manganous  salt  is  enough  to  so  affect  the  conditions  of 
equilibrium  that  titrations  in  moderate  volumes  (100  cm3  to 
500  cm3)  and  in  presence  of  hydrochloric  acid  (5  cm3  to  15  cm3 
of  the  strong  acid)  may  be  conducted  with  safety  and  reasonable 
rapidity,  either  with  or  without  sulphuric  acid,  at  the  ordinary 
atmospheric  temperature. 

Volume  at  Beginning  of  Titration  =  140  cm8. 
Temperature  =  20°-24°  C. 


Ammonium 
oxalate. 

HCl. 

strongest. 

MnCl2. 
4H20. 

BaClo. 

SrCl,. 

CaCla- 

MgCl,. 

KMn04 
used. 

Error. 

cm8 
25 

cm8 

5 

grm. 

0.5 

grm. 

grm. 

grm. 

grm. 

C1113 

26.05 

cm3 
0.00 

25 
25 
25 
25 
25 

5 
5 
5 
5 
5 

• 

¥ 

¥ 

¥ 

¥ 

27.45 
26.50 
26.53 
26.36 
26.13 

1.40+ 
0.45+ 
0.48+ 
0.35+ 
0.08+ 

25 
25 
25 
25 

5 
5 
6 
5 

0.5 
0.5 
0.5 
0.5 

2 

¥ 

¥ 

¥ 

26.05 
26.10 
26.10 
26.05 

0.00 
0.05+ 
0.05+ 
0.00 

25 
25 
25 
25 

10 
10 
10 
10 

1.0 
1.0 
1.0 
1.0 

2 

V  .. 

V 

V 

26.10 
26.05 
26.06 
26.11 

+0.05 
0.00 
0.01+ 
0.06+ 

XXVIII 

THE  ESTIMATION  OF  IKON  IN  THE  FERRIC  STATE 
BY  REDUCTION  WITH  SODIUM  THIOSULPHATE 
AND  TITRATION  WITH  IODINE. 

BY  JOHN  T.  NORTON,  JB.* 

THE  action  of  sodium  thiosulphate  on  ferric  iron  has  long 
been  known  and  depends  upon  the  following  reaction: 

2FeCl3  +  2Na2S2O8  =  SFeCL,  +  Na2S4O6  +  2NaCl. 

As  early  as  1859  Schererf  proposed  a  method  for  the 
estimation  of  ferric  iron  depending  on  the  above  reaction. 
Scherer's  method  of  procedure  was  to  act  upon  a  solution  of 
ferric  chloride  with  sodium  thiosulphate  until  the  purple 
color  produced  by  the  interaction  of  these  two  salts  just 
vanished.  Mohr's  $  experimental  tests  of  this  process  were 
not  successful.  A  year  or  two  later  Kremer  and  Landolt,§ 
after  a  careful  investigation  of  Scherer's  process,  recommended 
it  with  the  modification  that  any  free  hydrochloric  acid  pres- 
ent should  be  neutralized  by  sodium  acetate  until  the  solution 
assumed  a  red  color,  just  enough  hydrochloric  acid  added  to 
destroy  this  red  color,  and  sodium  thiosulphate  run  into  the 
solution  in  slight  excess.  When  the  liquid  became  perfectly 
colorless  and  gave  no  reaction  for  ferric  iron  with  potassium 
sulphocyanide,  the  excess  of  sodium  thiosulphate  was  titrated 
back  with  iodine  and  starch.  The  authors  also  state  that 
the  ferric  iron  should  not  be  present  in  concentrated  solution. 
Very  good  results  were  claimed  for  this  process,  but  it 
apparently  gained  but  slight  recognition. 

*  From  Am.  Jour.  Sci.,  viii,  25. 

t  Gel.  Anzeig.  k.  Bayrisch.  Acad.,  Aug.  31, 1859. 

$  Ann.  Chem.  Pharm.,  cxiii,  260.  §  Zeitschr.  anal.  Chem.,  i,  214. 


ESTIMATION  OF  IRON  IN  THE  FERRIC  STATE.     231 

Oudemanns,*  who  was  the  next  to  study  the  action  of  ferric 
iron  and  sodium  thiosulphate,  claimed  that  the  addition  of  a 
small  quantity  of  cupric  salt  to  the  iron  solution  hastened  the 
reducing  action  of  the  sodium  thiosulphate.  Mohr,f  however, 
condemned  this  method  also  as  unreliable,  both  because  the 
sodium  thiosulphate  acted  upon  the  copper  as  well  as  the  iron 
and  also  because  the  potassium  sulphocyanide,  added  as  an 
indicator  of  the  completeness  of  the  reduction,  produced  a 
precipitate  of  cupric  sulphocyanide  which  interfered  with 
the  reaction.  In  a  second  paper  Oudemanns  J  reiterated 
his  former  statement  as  to  the  accuracy  of  his  method  but 
advised  the  use  of  a  smaller  quantity  of  the  cupric  salt. 
An  improvement  on  Oudemanns'  process  was  proposed  by 
Haswell,§  who  mixed  the  moderately  acid  solution  of  ferric 
chloride  in  the  presence  of  a  cupric  salt  with  a  few  drops  of 
sodium  salicylate  and  then  reduced  with  sodium  thiosulphate 
previously  standardized  upon  a  known  quantity  of  iron  by 
the  same  process  and  estimated  the  excess  by  potassium 
dichromate.  Bruel  ||  modified  this  process  by  operating 
without  the  copper  solution,  relying  merely  on  the  discharge 
of  the  violet  color  in  a  boiling  solution  by  sodium  thiosul- 
phate standardized  on  a  ferric  solution  of  known  strength. 

Although  considerable  work  has  been  done  on  the  reaction 
between  ferric  iron  and  sodium  thiosulphate,  no  process 
depending  upon  this  reaction  has  obtained  acceptance.  In 
view,  therefore,  of  previous  work  on  the  action  of  hydro- 
chloric acid  upon  sodium  thiosulphate  If  and  with  the  idea  that 
a  careful  control  of  the  dilution  and  quantity  of  acid  present 
might  greatly  better  the  accuracy  of  the  method,  it  has 
seemed  to  me  to  be  desirable  to  study  this  process  again  in 
detail. 

The  ferric  oxide  employed  in  the  experiments  was  prepared 
with  great  care  by  the  ignition  of  ferrous  oxalate  obtained  by 

*  Zeitschr.  anal.  Chem.,  vi,  129.         t  Titrirmethode,  5et  Aufl.,  294. 
J  Zeitschr.  anal.  Chem.,  ix,  362. 
§  Kepertorium  der  analytischen  Chem.,  i,  179. 
||  Compt.  rend.,  xcvii,  954. 
1  Am.  Jour.  Sci.,  vii,  287.    This  volume,  p.  206. 


232    ESTIMATION  OF  IRON  IN  THE  FERRIC   STATE 

acting  with  oxalic  acid  on  pure  ammonium  ferrous  sulphate. 
To  ascertain,  however,  if  this  oxide  contained  any  impurity, 
about  0.5  of  a  grm.  was  put  into  a  porcelain  boat  and  sub- 
mitted to  the  action  of  a  current  of  hydrochloric  acid  gas  and 
chlorine  at  a  temperature  of  about  280°  C.  (according  to  a 
process  recently  described  from  this  laboratory  *)  until  all  the 
ferric  salt  is  volatilized  hi  the  form  of  ferric  chloride.  A 
residue  of  0.0010  grm.  for  every  0.5  of  a  grm.  of  the  oxide 
was  found,  and  this  correction,  small  for  the  amounts 
generally  used,  has  been  applied  in  the  following  determina- 
tions. The  sodium  thiosulphate  used  was  taken  in  nearly  ^ 
solution  and  was  standardized  against  an  approximately 
decinormal  solution  of  iodine  which  had  been  determined  by 
comparison  with  decinormal  arsenious  acid  made  from  care- 
fully resublimed  arsenious  oxide. 

In  those  experiments  which  deal  with  amounts  of  ferric 
oxide  not  exceeding  0.2  of  a  grm.,  measured  portions  of  a 
solution  of  ferric  chloride  made  of  known  strength  by  dis- 
solving about  2  grms.  of  the  pure  carefully  weighed  ferric 
oxide  in  20  cm3  of  strong  hydrochloric  acid  and  diluting  to 
one  liter,  were  drawn  from  a  burette.  In  the  case  of  the 
krger  quantities  of  ferric  oxide  the  salt  was  weighed  out, 
dissolved  in  hydrochloric  acid  and  brought  to  the  required 
dilution.  The  ferric  chloride,  either  drawn  from  the  burette 
or  prepared  directly  from  the  weighed  oxide,  was  diluted  with 
water,  a  drop  of  potassium  sulphocyanide  added  to  serve  as 
an  indicator  and  an  excess  of  sodium  thiosulphate  was  run 
in  until,  after  standing  for  a  few  minutes,  the  solution  became 
perfectly  colorless,  and  the  excess  of  sodium  thiosulphate 
was  then  titrated  back  with  decinormal  iodine  after  the 
addition  of  starch. 

Several  sources  of  error  are,  plainly,  possible  in  the  process : 
incompleteness  in  the  reduction  of  the  ferric  salt ;  decomposition 
of  the  thiosulphate  by  the  acid,  resulting  in  the  subsequent 
over-run  of  iodine;  the  possible  tendency  of  the  ferric  salt 
under  concentration  to  oxide  the  thiosulphate  to  the  condition 
*  Gooch  and  Havens,  Am.  Jour.  Sci.,  vii,  370.  This  volume,  p.  215. 


BY  DEDUCTION  WITH  SODIUM  THIOSULPHATE.   233 


of  the  sulphate  rather  than  to  that  of  the  tetrathionate ;  and 
finally  the  oxidizing  action  of  the  air,  which  may  tend  to 
keep  up  progressive  oxidation  of  the  iron  salt  and  excessive 
expenditure  of  thiosulphate.  The  first  three  sources  of  diffi- 
culty tend  to  produce  errors  of  deficiency;  the  fourth  an 
error  of  excess. 

The  first  step  in  the  experimental  study  of  the  process  was 
to  determine  the  effect  of  varying  dilution  upon  the  estimation 
of  a  given  quantity  of  iron  reduced  by  sodium  thiosulphate, 
taken  in  practically  uniform  excess  above  the  amount  theo- 
retically required,  in  the  presence  of  1  cm3  of  hydrochloric 
acid. 

TABLE  I. 


Exp. 

Fe203 
taken. 

Fe203 
corrected. 

Dilution. 

HC1. 

NaAO, 

in  excess. 

Fe208 
found. 

Error. 

grm. 

grm. 

cm* 

cm8 

cm» 

grm. 

grm. 

(1) 

0.1000 

0.0998 

100 

1 

18.08 

0.0957 

0.0041- 

2 

0.1000 

0.0998 

200 

20 

0.0966 

0.0032- 

(3) 

0.1000 

0.0998 

300 

17.56 

0.0995 

0.0003- 

(4) 

0.1000 

0.0998 

400 

17.16 

0.0998 

0.0000 

M 

0.1000 

0.0998 

600 

17.76 

0.0996 

0.0002- 

(6) 

0.1000 

0.0998 

800 

17.65 

0.0993 

0.0005- 

(7) 

0.1000 

0.0998 

1000 

18.02 

0.0988 

0.0010- 

(8) 

0.1000 

0.0998 

1200 

17.95 

0.0977 

0.0021- 

(9) 

0.1000 

0.0998 

1400 

17.99 

0.0965 

0.0033- 

(10) 

0.1000 

0.0998 

1600 

1 

18.01 

0.0947 

0.0051- 

(11) 

0.2001 

0.1997 

400 

2 

27.05 

0.2029 

0.0032+ 

(12) 

0.2001 

0.1997 

800 

2 

15.95 

0.1998 

0.0001+ 

(13) 

0.4998 

0.4988 

1000 

2 

22.36 

0.5104 

0.0116+ 

(14) 

0.5051 

0.5041 

1800 

4 

15.27 

0.5026 

0.0015— 

(15) 

0.4002 

0.-3994 

1500 

4 

27.29 

0.3996 

0.0002+ 

(16) 

0.7502 

0.7487 

1000 

1 

9.73 

0.7572 

0.0085+ 

(17) 

0.7029 

0.7015 

2000 

4 

12.67 

0.7004 

0.0011- 

This  table  shows  plainly  that  with  quantities  of  ferric  oxide 
present  up  to  0.1  grm.  the  dilution  can  vary  from  400  cm3  to 
1000  cm3  for  each  cm3  of  strong  hydrochloric  acid  and  still 
give  excellent  results.  At  a  dilution  greater  than  1000  cm3 
the  action  of  the  thiosulphate  is  evidently  incomplete,  and  at 
a  smaller  dilution  than  400  cm3  the  decomposing  action  of  the 
acid  on  the  thiosulphate  becomes  noticeable.  When  larger 
quantities  of  iron  oxide  are  dealt  with,  it  appears  that  the 


234    ESTIMATION  OF  IRON  IN  THE  FERRIC  STATE 


dilution  ought  to  be  increased  proportionally  with  the  quantity 
of  ferric  oxide  present  as  well  as  with  that  of  the  acid.  This 
is  illustrated  in  experiments  9-15  of  the  table.  On  this 
account  it  seems  necessary,  assuming  that  the  quantity  of 
acid  present  is  always  kept  within  the  maximum  strength 
mentioned,  1  cm3  to  400  cm3,  to  regulate  the  dilution  from  the 
approximate  quantity  of  the  iron  so  that  not  less  than  400  cm3 
of  water  shall  be  used  to  every  0.1  grm.  of  iron  oxide  present. 
Under  properly  regulated  conditions  of  dilution  as  regards 
acid  and  the  iron  salt,  the  reduction  is  completed  in  from  five 
to  ten  minutes. 

Great  excesses  of  acid,  however,  contrary  to  the  statement 
of  Kremer,*  retard  the  reduction  greatly,  and,  hi  spite  of  the 
tendency  of  the  thiosulphate  to  decomposition  and  the  pro- 
duction of  errors  of  deficiency  under  such  circumstances,  plus 
errors  due  to  partial  oxidation  come  to  light.  This  fact  appears 
in  the  following  table,  which  records  the  results  of  processes 
lasting  many  hours. 

TABLE  II 


Exp. 

Fe20s 
taken. 

Fe208 
corrected. 

DUution. 

HC1. 

Na2S20, 
in  excess. 

Fe208 
found. 

Error. 

grm. 

grm. 

cm8 

cm* 

cm3 

grm. 

grm. 

(18) 

0.5012 

0.6002 

1700 

10 

25.99 

5308 

0.0306+ 

(19) 

0.7512 

0.7497 

1200 

16 

57.8 

7685 

0.0188+ 

(20) 

0.7520 

0.7505 

2000 

16 

56.4 

7983 

0.0478+ 

(21) 

0.7520 

0.7505 

1700 

15 

27.2 

7627 

0.0122+ 

As  to  the  temperature  at  which  the  reduction  should  be 
made,  my  experience,  contrary  to  that  of  Kremer,  goes  to  show 
that  no  elevation  above  atmospheric  conditions  is  necessary ; 
under  the  conditions  of  acidity  and  dilution  laid  down,  the 
process  of  reduction  is  complete  within  ten  minutes  after  the 
introduction  of  the  thiosulphate ;  moreover,  former  experience  f 
shows  clearly  the  danger  of  submitting  mixtures  of  sodium 
thiosulphate  and  acid  to  temperatures  much  above  the  ordinary. 
On  the  other  hand,  artificial  reduction  of  temperature  tends 

*  Zeitschr.  anal.  Chem.,  i,  214. 

t  Am.  Jour.  Sci.,  vol.  vii,  287.    This  volume,  p.  206. 


EY  REDUCTION  WITH  SODIUM  THIOSULPHATE.     235 


to  retard  the  action  to  an  impossible  degree.  Thus,  in  an 
experiment  it  took  five  minutes  to  reduce  0.0500  of  ferric 
oxide  at  21  °  C.  completely  at  a  dilution  of  200  cm3  and  in 
the  presence  of  |-  cm3  of  hydrochloric  acid ;  under  conditions 
otherwise  precisely  similar  excepting  that  the  temperature  was 
lowered  to  0°  C.,  the  action  lingered  forty-five  minutes. 

Lastly,  the  question  as  to  the  excess  of  thiosulphate  necessary 
to  complete  the  reduction  within  a  reasonable  time  must  be 
considered.  In  nearly  all  previously  recorded  experiments 
the  excess  of  thiosulphate  was  not  less  than  15  cm3  of  the 
^  solution.  The  following  table  shows  the  effect  of  dimin- 
ishing this  excess. 

TABLE  HI. 


Kxp. 

Fe,08 
taken. 

Fe,0s 
corrected. 

Dilution. 

HC1. 

NaAO, 

in  excess. 

aa 

Error. 

grin* 

grm. 

cm3 

cm3 

cms 

grm. 

grm. 

(22) 

0.0250 

0.0250 

400 

12.2 

0.0241 

0.0009- 

(23) 

0.0500 

0.0499 

400 

12.2 

0.0495 

0.0004- 

(24) 

0.0500 

0.0499 

400 

13.66 

0.0493 

0.0006- 

(25) 

0.1000 

0.0998 

400 

1 

7.31 

0.0984 

0.0014- 

(26) 

0.1000 

0.0998 

400 

1 

7.63 

0.0972 

0.0026- 

(27) 

0.1001 

0.0999 

400 

1 

12.88 

0.1007 

0.0008+ 

(28) 

0.1498 

0.1495 

600 

1* 

11.97 

0.1475 

0.0020- 

(29) 

0.1996 

0.1992 

800 

2 

12.43 

0.1980 

0.0012- 

From  the  above  experiments  taken  in  connection  with  those 
of  Table  I  it  is  clear  that  there  should  always  be  present  an 
excess  of  at  least  15  cm3  of  the  T^  solution  of  sodium  thiosul- 
phate. If  the  quantity  of  hydrochloric  acid  is  kept  very  low 
there  is  no  reason  why  this  excess  of  thiosulphate  could  not 
be  considerable  without  producing  any  disturbing  effect. 
Practically,  however,  the  presence  of  an  excess  between  the 
limits  of  15  cm8  and  35  cm3  of  the  ^  solution  has  been  found 
to  give  the  most  satisfactory  results. 

To  recapitulate,  then,  it  has  been  shown  that  the  dilution 
must  be  at  least  400  cm3  for  each  0.1  of  a  grm.  of  iron  oxide 
present,  that  the  quantity  of  acid  should  never  exceed  1  cm3 
of  the  strong  acid  to  each  400  cm3  of  water,  that  the  time  of 
reduction  must  be  short  to  avoid  progressive  oxidation,  that 


236    ESTIMATION  OF  IRON  IN  THE  FERRIC  STATE 


the  temperature  of  the  solution  should  be  kept  at  the  normal 
temperature  of  the  atmosphere,  and  finally  that  the  excess  of 
sodium  thiosulphate  present  should  never  be  less  than  15  cm3 
of  the  £f  solution.  In  the  case  of  large  dilution  the  use  of 
freshly  boiled  water  is  recommended  so  as  to  avoid  the 
reoxidizing  effect  of  the  air  upon  the  reduced  iron.  In  the 
experiments  included  in  the  following  table,  the  above 
precautions  were  closely  adhered  to  and  manifestly  satisfactory 
results  were  obtained. 

TABLE  IV. 


Exp. 

Fe203 
taken. 

Fe203 
corrected. 

Dilution. 

HC 

i. 

Excess 
Na,S203 

found. 

Error. 

grin. 

grm. 

cm8 

cm 

3 

cm* 

grm. 

grm. 

(30) 

0.0125 

0.0125 

200 

23.5 

0.0125 

0.0000 

(31) 

0.0250 

0.0250 

400 

21.98 

0.0250 

0.0000 

(32) 

0.0250 

0.0250 

400 

17 

0.0250 

0.0000 

(33) 

0.0250 

0.0250 

400 

17 

0.0250 

0.0000 

(34) 

0.0500 

0.0499 

400 

24 

0.0498 

0.0001- 

(35) 

0.0500 

0.0499 

400 

19 

0.0498 

0.0001- 

(36) 

0.0500 

0.0499 

400 

15.1 

0.0497 

0.0002- 

(37) 

0.0500 

0.0499 

400 

19 

0.0498 

0.0001- 

(38) 

0.1001 

0.0999 

400 

23.1 

0.0993 

0.0006- 

(39) 

0.1001 

0.0999 

400 

17.93 

0.0997 

0.0002— 

(40) 

0.1001 

0.0999 

400 

22.92 

0.0997 

0.0002- 

(41) 

0.1001 

0.0999 

400 

18 

0.0997 

0.0002- 

(42) 

0.1001 

0.0999 

400 

16 

0.0996 

0.0003- 

(43) 

0.1498 

0.1495 

600 

j 

- 

23.26 

0.1493 

0.0002- 

(44) 

0.1498 

0.1495 

600 

^ 

- 

16.66 

0.1493 

0.0002- 

(45) 

0.1498 

0.1495 

600 

- 

26.87 

0.1475 

0.0020- 

(46) 

0.1996 

0.1992 

800 

2 

22.38 

0.1990 

0.0002- 

(47) 

0.1996 

0.1992 

800 

2 

17.29 

0.1999 

0.0007+ 

(48) 

0.1996 

0.1992 

800 

2 

22.20 

0.1991 

0.0001- 

(49) 

0.4045 

0.4037 

1600 

4 

16.03 

0.4042 

0.0005+ 

(50) 

0.4045 

0.4037 

1600 

4 

16.2 

0.4023 

0.0014- 

(51) 

0.4018 

0.4010 

1600 

4 

16.34 

0.4007 

0.0003- 

(52) 

0.5051 

0.5041 

1800 

4 

15.27 

0.5026 

0.0015- 

As  seen  in  the  table  this  process  is  very  accurate,  especially 
in  the  use  of  small  amounts  of  ferric  oxide.  The  introduc- 
tion of  cupric  sulphate  as  recommended  by  Oudemanns,  or  of 
sodium  salicylate  according  to  Haswell's  method,  seems  to  be 
unnecessary  and  only  complicates  the  process. 

In  treating  ferric  oxide,  the  following  method  of  procedure 
is  recommended.  Dissolve  an  amount  not  exceeding  0.2  grm. 
of  the  oxide  in  hydrochloric  acid,  evaporate  to  a  pasty  mass 


BY  REDUCTION   WITH  SODIUM  THIOSULPHATE.     237 

dilute  to  about  800  cm3  with  freshly  boiled  water,  add  a  drop  of 
potassium  sulphocyanide,  and  into  this  solution  run  50  cm3  of 
approximately  ^  sodium  thiosulphate;  allow  the  liquid  to 
stand  until  perfectly  colorless  and  determine  the  excess  of 
thiosulphate  by  ^  iodine  and  starch.  For  quantities  of  iron 
oxide  up  to  0.2  of  a  gram  this  process  is  quick  and  most  accu- 
rate; when  care  is  taken  to  preserve  the  relations  of  acidity 
and  dilution,  twice  the  amount  of  ferric  oxide  mentioned 
above  may  be  handled. 


XXIX 

THE  DETERMINATION   OF  TELLUROUS   ACID  IN 
PRESENCE   OF  HALOID  SALTS. 

BY  F.  A.  GOOCH  AND  C.  A.  PETEKS.* 

THE  estimation  of  tellurous  acid  by  oxidation  with  excess  of 
potassium  permanganate  (either  in  acid  or  alkaline  solution), 
destruction  of  the  higher  oxides  of  manganese  or  the  manga- 
nate  by  standard  oxalic  acid  in  presence  of  sulphuric  acid,  and 
titration  of  the  residual  oxalic  acid  by  more  permanganate,  has 
been  shown  by  Braunerf  to  be  feasible.  The  tendency  of  the 
permanganate  to  throw  off  too  much  oxygen  when  the  oxida- 
tion is  made  in  solutions  strongly  acidified  with  sulphuric  acid 
(as  must  be  the  case  if  the  tellurous  oxide  is  to  be  held  perma- 
nently in  solution  by  sulphuric  acid)  necessitates  the  applica- 
tion of  a  considerable  correction.:]:  Fortunately,  however,  as 
has  been  shown  in  a  former  paper  from  this  laboratory,!  when 
the  tellurous  oxide  is  dissolved  originally  in  an  alkaline  hydrox- 
ide and  the  solution  made  acid  only  to  a  limited  degree  with 
sulphuric  acid  either  before  or  after  oxidation  by  the  perman- 
ganate, no  correction  appears  to  be  necessary.  Thus,  when 
an  excess  of  permanganate  is  added  to  the  alkaline  solution, 
followed  by  an  excess  of  oxalic  acid  and  sulphuric  acid  to  an 
amount  not  exceeding  5  cm3  of  the  [1 : 1]  mixture  with  water, 
the  titration  of  the  residual  oxalic  acid  by  more  permanganate 
(after  heating  to  80°  C.)  leads  to  results  which  give  no  indi- 
cation of  over-decomposition  of  the  permanganate;  so  also, 
when  the  process  is  similarly  conducted  excepting  that  before 
addition  of  the  permanganate  the  original  alkaline  solution  is 
acidified  with  sulphuric  acid  [1 : 1]  to  an  amount  1  cm3  in  excess 

*  From  Am.  Jour.  Sci.,  viii,  122.  t  Jour.  Chem.  Soc.,  lix,  238. 

J  Loc.  cit,  p.  249. 

§  Gooch  and  Danner,  Am.  Jour.  Sci.,  xliv,  301.    Volume  I,  p.  145. 


DETERMINATION  OF  TELLUROUS  ACID,  ETC.       239 

of  that  necessary  to  redissolve  the  first  precipitate,  the  results 
are  theoretically  accurate,  and  in  close  agreement  with  those 
obtained  by  the  former  procedure. 

In  the  presence  of  free  hydrochloric  acid  the  action  of  the 
permanganate  upon  tellurous  acid  has  been  shown  by  Brauner* 
to  be  irregular  and  excessive,  and  the  irregularity  cannot  be 
corrected  (as  in  the  titration  of  ferrous  salts  in  presence  of 
hydrochloric  acid)  by  the  addition  of  a  manganous  salt  accord- 
ing to  the  well-known  procedure  of  Kesslerf  and  Zimmer- 
mann4  So  far  as  appears,  however,  there  should  be  nothing 
to  prevent  the  accurate  determination  of  tellurium  in  tellurous 
compounds  in  the  presence  of  chlorides  by  the  permanganate 
process  providing  the  first  oxidation  is  made  in  alkaline  solu- 
tion, and  the  second  oxidation  carried  out  with  such  precau- 
tions as  are  necessary  to  a  correct  determination  of  oxalic  acid 
by  permanganate  hi  presence  of  hydrochloric  acid ;  for  the 
special  danger  of  over-action  on  the  part  of  the  permanganate 
cannot  exist  while  the  solution  is  alkaline,  and  has  passed 
when  the  tellurite  has  become  a  tellurate  and  before  the  solu- 
tion is  made  acid.  As  to  the  proper  conditions  for  the  titra- 
tion of  oxalic  acid  by  permanganate  we  have  shown  recently§ 
that  the  presence  of  a  manganous  salt  is  necessary  and  suf- 
cient  to  secure  regularity  of  action  when  a  considerable 
amount  of  hydrochloric  acid  is  in  the  solution;  when  the 
amount  is  small  —  so  much  as  would  be  formed  in  the  decom- 
position of  a  gram  or  two  of  halogen  salt  of  tellurium  —  the 
disturbing  effect  under  ordinary  conditions  of  work  is  prob- 
ably inappreciable,  but  even  in  such  a  case  it  is  better  to  work 
in  the  presence  of  a  manganous  salt  for  the  reason  that  the 
titration  of  the  oxalic  acid  may  then  be  made  at  the  ordinary 
atmospheric  temperature. 

In  the  following  table  are  gathered  the  results  of  experi- 
ments made  with,  and  without,  the  addition  of  the  manganous 
salt. 

*  Loc.  cit.,  p.  241.  t  Ann.  Phys.  cxviii,  48 ;  cxix,  225,  226. 

|  Ann.  Chera.  (Liebig),  ccxiii,  302. 

§  Am.  Jour.  Sci.,  vii,  p.  461.    This  volume,  p.  222. 


240 


DETERMINATION  OF  TELLUROUS  ACID 


TABLE  I. 

O  =  16,  Te  =  127.5. 

Volume  at  beginning,  150  cm8. 

Temperature  of  titration,  60-80°  C. 


TeO, 

taken. 

NaCl. 

5f2- 

MnCl2.4H20. 

TeOg 

found. 

Error. 

grm. 

0.1000 
0.1000 
0.1000 
0.1000 
0.0650 

grm. 

0.4 
0.4 
0.4 
1.0 
1.0 

cm8 
5 
5 
5 
5 
5 

grill. 

grm. 

0.1003 
0.1000 
0.1004 
0.1003 
0.0653 

grm. 

0.0003+ 
0.0000 
0.0004+ 
0.0003+ 
0.0003+ 

B. 

Temperature  of  titration,  20-26°  C. 

0.0700 
0.0700 
0.0700 
0.1000 

0.4 
0.4 
0.4 
0.4 

5.7 
5.7 
5.7 
5.7 

1.0 
1.0 
0.5 

0.5 

0.0705 
0.0698 
0.0701 
0.1008 

0.0005+ 
0.0002- 
0.0001+ 
0.0008+ 

The  tellurium  dioxide,  made  by  the  careful  ignition  of  the 
crystallized  basic  nitrate  obtained  by  oxidizing  tellurium  with 
nitric  acid,  was  dissolved  in  a  small  amount  of  sodium  hydrox- 
ide, the  halogen  salt  was  added  to  the  amount  shown,  the  per- 
manganate standardized  against  ammonium  oxalate  was  run 
in  until  its  characteristic  color  appeared,  standard  ammonium 
oxalate  was  added  in  excess  of  the  quantity  required  to  reduce 
the  excess  of  permanganate,  manganate,  and  higher  oxides, 
and  the  solution  was  heated  with  enough  sulphuric  acid  [1:1] 
to  neutralize  the  alkaline  hydroxide  and  have  an  excess  of 
about  5  cm3.  In  the  experiments  of  Section  A  the  liquid  was 
heated  to  60°  -80°  C.  to  dissolve  the  oxides  at  the  final  titra- 
tion begun  at  that  temperature ;  in  those  of  Section  B, 
manganous  chloride  (0.5  to  1  gram)  was  added,  so  that  the 
reduction  of  the  higher  oxides  of  manganese  and  the  final 
titration  of  the  excess  of  oxalic  acid  might  take  place  at  the 
ordinary  temperature  of  the  room. 

Plainly  the  presence  of  the  chloride  does  not  interfere 
materially  in  the  determination  of  the  tellurium  by  this 
process  whether  the  titration  is  made  at  a  high  or  low 
temperature. 


IN  PRESENCE  OF  HALOID  SALTS. 


241 


It  appears,  also,  upon  putting  the  matter  to  the  test,  that 
fairly  good  determinations  of  tellurous  acid  may  be  made 
similarly  in  the  presence  of  a  bromide,  provided  the  titration 
is  made  at  the  atmospheric  temperature  in  the  presence  of  a 
sufficiency  (0.5  gram  to  1  gram)  of  a  manganous  salt  and  of 
an  excess  of  sulphuric  acid  limited  to  about  5  cm3  or  less  of  the 
12.5  per  cent  mixture.  At  the  higher  temperatures  bromine 
is  liberated  at  once  from  the  acid  solution  by  the  permanganate. 
The  experimental  results  are  given  in  Table  II. 

TABLE   II. 

0  =  16,  Te  =  127.5. 

Volume  at  beginning,  150  cm8. 

Temperature  of  titration,  24°-26°  C. 


Te02 
taken. 

NaCl 

KBr. 

Sfc 

MnCl2. 
4H2O. 

TeO, 
found. 

Error. 

grm. 

grm. 

grm. 

cms 

grm. 

gnn. 

grm. 

0.1000 

.  . 

0.5 

20 

1.0 

0.1022 

0.0022+ 

0.3000 

.  . 

1.5 

25 

1.0 

0.3030 

0.0030+ 

0.0650 

0.5 

1 

•  1.0 

O.OG61 

0.0011+ 

0.0650 

0.5 

1 

1.0 

0.0647 

0.0003- 

0.1000 

0.5 

1 

1.0 

0.1002 

0.0002+ 

0.3000 

.  . 

0.5 

5 

0.5 

0.3010 

0.0010+ 

0.0650 

0.5 

0.5 

1 

1.0 

0.0661 

0.0011+ 

It  is  obvious,  therefore,  that  tellurous  acid  may  be  deter- 
mined with  a  fair  degree  of  accuracy  by  the  permanganate 
method  in  the  presence  of  chlorides  and  bromides,  provided 
the  first  oxidation  is  made  in  alkaline  solution  and  the  final 
titration  of  the  residual  oxalic  acid  is  made  at  ordinary 
temperatures  in  the  presence  of  a  manganous  salt  and  restricted 
amounts  of  free  sulphuric  acid. 

In  the  presence  of  an  iodide,  however,  the  case  is  different. 
Upon  acidifying  the  mixture  of  iodide  and  the  higher  oxygen 
compounds  of  manganese,  produced  in  the  action  of  the 
permanganate  upon  the  solution,  iodine  is  at  once  set  free,  and 
oxalic  acid  does  not  suffice  to  reconvert  it.  In  the  presence  of 
an  excess  of  potassium  iodide  the  higher  manganic  compounds 
are  completely  reduced  with  rapidity  and  the  iodine  liberated 
is  the  measure  of  the  excess  of  permanganate  over  that 

VOL.   II.  — 16 


242  DETERMINATION  OF  TELLUROUS  ACID 

required  to  oxidize  the  tellurous  acid ;  the  difference  between 
the  amount  of  permanganate  thus  indicated  and  that  originally 
introduced  should  determine  the  amount  of  the  tellurous  acid. 
It  is  upon  this  basis  that  Norris  and  Fay  *  have  founded  their 
excellent  iodometric  determination  of  tellurous  acid.  This 
process  consists  in  treating  the  alkaline  solution  of  tellurous 
acid  with  standard  permanganate  until  the  meniscus  of  the 
liquid  shows  a  deep  pink  color,  then  diluting  the  solution  with 
ice-water,  adding  potassium  iodide  and  sulphuric  acid,  and 
titrating  with  sodium  thiosulphate.  The  results  are  excellent. 

It  is  plain  that  any  agent  capable  of  converting  the  iodine 
to  hydriodic  acid  without  at  the  same  time  reducing  telluric 
acid  should  be  capable  of  measuring  the  excess  of  the 
permanganate,  and  so  the  amount  of  tellurous  acid  originally 
present.  We  find  that  the  standard  arsenite  made,  as  usual, 
by  dissolving  4.95  grams  of  pure  resublimed  arsenious  oxide 
to  the  liter  of  water  containing  potassium  bicarbonate  answers 
the  purpose  admirably,  and  possesses  the  further  advantage  of 
fixing  at  once  the  entire  standard  of  the  process,  the  strength 
of  the  permanganate  (approximately  -^  being  determined  by 
running  a  definite  volume  of  its  solution  into  water  containing 
potassium  iodide  (1  gram)  with  2  to  3  cm3  of  dilute  sulphuric 
acid  and  titrating  by  the  standard  arsenite  the  iodine  (set 
free  by  the  action  of  the  excess  of  permanganate  and  higher 
oxides)  after  neutralization  with  acid  potassium  bicarbonate. 
In  this  titration  of  iodine  by  the  arsenite  we  find  it  best  to 
dispense  with  the  starch  solution  usually  employed  to  secure 
the  end  reaction.  The  color  of  the  free  iodine  itself  is 
sufficiently  definite,  even  at  a  dilution  so  much  as  300  cm8, 
and  its  disappearance  under  the  action  of  the  arsenite  is  much 
sharper  than  that  of  the  blue  starch  iodide. 

In  Table  III  are  recorded  results  obtained  by  adding  the 
alkaline  solution  of  tellurous  oxide  to  100  cm8  of  water 
containing  0.5  gram  or  1  gram  of  potassium  iodide,  introducing 
the  standardized  potassium  permanganate  until  the  green  color 
of  the  manganate  appears  (about  30  cm3  of  the  ^  solution  for 

*  Am.  Chem.  Jour.,  xx,  278. 


IN  PRESENCE  OF  HALOID  SALTS. 


243 


every  0.1  gram  of  TeO2),  adding  a  few  cubic  centimeters  of 
dilute  sulphuric  acid,  followed,  when  the  solution  has  cleared 
and  separated  iodine,  by  an  excess  of  acid  potassium  carbonate, 
and  titrating  to  the  destruction  of  color  with  the  standard 
solution  of  arsenic.  It  is  essential,  in  order  that  oxygen  may 
not  go  to  waste  in  the  breaking  down  of  the  oxides,  that  more 
than  enough  iodide  should  be  present  when  the  solution  is 
acidified  to  complete  the  reduction  of  the  manganese  oxides, 
or  else,  that  the  arsenious  acid  should  be  present  in  suitable 
amount  before  the  sulphuric  acid  is  put  in.  This  latter 
procedure  may  be  used  in  case,  for  any  reason,  it  is  preferred 
not  to  introduce  more  iodide  into  the  solution  than  may  be 
present  originally :  when,  for  example,  a  direct  determination 
of  the  iodine  present  is  to  follow. 

TABLE  III. 
O  -  16,  Te  =  127.5. 


Te02 
taken. 

NaCL 

KBr. 

KI. 

Total 
volume 

NaOH 
present 
during 

TeO, 
found. 

Error. 

oxidation. 

griii* 

grin* 

grm. 

grm. 

cm3. 

grm. 

grm. 

grm. 

0.1000 

0.5 

160 

0.1 

0.1005 

0.0005+ 

0.1000 

0.5 

160 

0.1 

0.1001 

0.0001+ 

0.1000 

. 

0.5 

160 

0.1 

0.1003 

0.0003+ 

0.1000 

1.0 

250 

0.1 

0.1007 

0.0007+ 

0.2000 

1.0 

250 

0.2 

0.1997 

0.0003+ 

0.1000 

0.6 

0.5 

0.5 

250 

0.1 

0.1000 

0.0000 

0.2100 

1.0 

1.0 

1.0 

225 

0.2 

0.2105 

0.0005+ 

0.1000 

§ 

0.5 

160 

1.0 

0.1011 

0.0011+ 

0.2000 

1.0 

300 

2.0 

0.2009 

0.0009+ 

These  results  are  reasonably  good.  Like  those  of  Table  I 
they  would  be  brought  practically  in  the  average  to  the  figure 
demanded  by  theory  if  the  value  of  the  Committee  of  the 
German  Chemical  Society,  Te  =  127,  were  to  be  taken  instead 
of  Te  =  127.5,  the  value  of  Clarke  and  of  Richards. 


XXX 

AN    IODOMETRIC    METHOD    FOR    THE    ESTIMA- 
TION OF  BORIC  ACID. 

BY  LOUIS  CLEVELAND  JONES.* 

IN  a  recent  article,!  I  have  described  a  process  for  the  alka- 
limetric  estimation  of  boric  acid,  depending  upon  the  forma- 
tion of  a  strongly  acidic  compound  when  boric  acid  and  a 
polyatomic  alcohol  are  placed  together  in  solution.  The 
method  in  brief  consists  in  destroying  the  free  mineral  acid 
in  a  solution  containing  boric  acid,  by  means  of  a  mixture  of 
potassium  iodide  and  iodate,  bleaching  the  liberated  iodine  by 
sodium  thiosulphate,  adding  the  indicator  phenolphthalein 
and  sufficient  standard  solution  of  caustic  soda  to  give  a 
faint  alkaline  coloration,  bleaching  by  a  small  amount  of 
mannite  and  adding  caustic  soda  again  to  alkalinity,  and 
thus  alternating  with  mannite  and  alkali  until  the  alkaline 
coloration  produced  is  permanent.  The  amount  of  sodium 
hydroxide  used  represents  the  amount  of  acidity  developed 
by  the  influence  of  the  mannite  upon  the  boric  acid  present, 
according  to  the  hypothesis  that  the  molecule  B2O8  acts  as 
two  molecules  of  a  univalent  acid,  HOBO. 

On  making  further  study  of  this  reaction,  I  have  found 
that  the  acid  developed  by  the  combination  of  boric  acid  and 
mannite  is,  under  certain  definite  conditions,  sufficiently 
strong  to  liberate,  quantitatively,  from  a  mixture  of  potassium 
iodide  and  iodate,  the  amount  of  iodine  required  on  the 
supposition  that  each  molecule  of  metaboric  acid  (HOBO) 
acts  in  a  manner  similar  to  a  univalent  mineral  acid  under 
the  same  conditions.  (5KI  +  KIO8  +  6HOBO  =  3I2  + 

*  From  Am.  Jour.  Sci.,  viii,  127. 

t  Am.  Jour.  Sci.,  vii,  147.    This  volume,  p.  182. 


THE  ESTIMATION  OF  BORIC  ACID.  245 

6KOBO  +  3H2O.)  Obviously,  this  reaction  depends  upon 
the  behavior  of  the  acidic  boromannite  compound  as  a  strong 
acid,  stronger  than  acetic,  tartaric,  or  citric  acid;  for  these 
acids  have  been  found  by  Furry  *  to  be  incapable  of  liberat- 
ing iodine  regularly  from  a  mixture  of  iodide  and  iodate. 
Conditions  which  tend  to  increase  the  acidic  activity  of  this 
compound  are  concentrated  solutions  and  moderately  low 
temperatures.f 

Glycerine  acts  in  general  like  mannite  to  produce  acidic 
compounds  with  boric  acid;  and  hi  a  preliminary  way,  the 
relative  acidity  of  the  products  formed  by  these  two  poly- 
atomic alcohols  with  boric  acid  may  be  indicated  by  the  results 
of  two  experiments  in  which  the  iodine  liberated  from  a 
mixture  of  potassium  iodide  and  iodate,  proportionately  to 
the  time  required  for  the  liberation,  is  taken  as  a  measure  of 
the  strengths  of  the  acids  developed. 

Equal  amounts  (10  cm3)  of  a  standard  solution  of  boric 
acid,  prepared  from  the  anhydride,  J  were  drawn  into  separate 
Erlenmeyer  flasks  and  a  neutral  solution  of  iodide  and  iodate 
added  to  each  in  an  amount  sufficient  to  liberate  iodine  in 
quantities  corresponding  to  the  acid  used.  One  solution 
was  treated  with  glycerine  enough  to  constitute  one-half  the 
entire  volume  of  the  liquid:  mannite  (about  5  grms.)  was 
added  to  the  other.  The  thiosulphate  required  immediately 
and  after  definite  periods  of  tune,  is  shown  for  each  solution 
in  the  following  table. 

The  solution  of  boric  acid  contained  7.706  grm.  per  liter. 
The  thiosulphate  was  0.0999  normal.  According  to  theory, 
the  amount  of  thiosulphate  required  for  10  cm3  of  the  boric 

*  Am.  Chem.  Jour.,  vi,  341. 

t  Magnanini,  Gaz.  chim.  Hal.  xx,  428,  xxi,  134 ;  and  Lambert,  Compt.  rend., 
CTiii,  1016, 1017. 

J  The  recrystallized  hydrous  boric  acid  should  be  fused  in  a  platinum  dish 
and,  after  cooling  and  breaking  into  small  pieces,  the  desired  amount  placed 
in  a  small  weighed  platinum  crucible  and  again  fused  until  no  more  water 
escapes.  After  cooling  and  weighing,  the  boric  oxide  may  be  separated  from 
the  crucible,  or  with  it  placed  in  warm  water,  dissolved  and  made  up  to  a 
definite  volume. 


246 


AN  IODOMETRIC  METHOD  FOR 


TABLE  I. 


Bj08  solution  (10  cm») 
with  maimite. 

Time. 

B20S  solution  (10  cm8) 
with  glycerine. 

Thiosulphate  required. 

Thiosulphate  required. 

cm* 

cm' 

18.60 
21.30 

Immediately. 
After  15  minutes. 

8.48 
10.50 

22.00 

After  30  minutes. 

11.15 

22.05 

After    2  hours. 

11.60 

acid  solution  is  22.02  cm3.  From  these  data  we  may  observe 
that  at  the  end  of  30  minutes,  in  the  solution  containing 
mannite,  practically  the  theoretical  amount  of  thiosulphate 
had  been  used,  while  only  about  50  per  cent  of  that  amount 
had  been  required  to  bleach  the  iodine  liberated  by  the 
glycerine  compound.  Obviously,  mannite  forms  with  boric 
acid  a  more  acidic  compound  than  glycerine  does,  and,  from 
the  indication  given  in  the  above  experiments,  may  be  relied 
upon,  under  certain  conditions,  to  liberate  the  theoretical 
amount  of  iodine.  If,  from  the  iodide  and  iodate  used  to 
destroy  the  excess  of  mineral  acid  already  present,  the  boric 
acid,  upon  the  addition  of  mannite  does  liberate  iodine  regu- 
larly—  as  the  previous  experiments  seem  to  indicate  —  this 
liberated  iodine  should  form  a  most  convenient  measure  of 
the  boric  acid  present. 

On  studying  the  conditions  requisite  for  the  complete 
liberation  of  iodine  according  to  theory,  several  important 
points  have  come  to  light. 

It  has  not  been  found  possible  under  any  conditions  to  rely 
upon  the  immediate  liberation  of  the  full  amount  of  iodine: 
a  certain  period  of  time  is  required  for  the  completion  of 
the  reaction.  When  the  solution  is  of  small  volume  and 
saturated  with  mannite,  the  reaction  goes  to  the  end  most 
quickly — sometimes  almost  immediately  —  but  there  is  this 
limitation,  which  must  be  made  emphatic,  viz.:  that  if  the 
solution  of  boric  acid  is  too  concentrated  —  near  saturation  — 
the  boric  acid  alone,  when  the  iodate  and  iodide  are  added  to 
destroy  any  other  free  acid  present,  throws  out  some  iodine 


THE  ESTIMATION  OF  BORIC  ACID.  247 

and  on  bleaching  with  thiosulphate  a  starting-point  is  ob- 
tained at  which  some  of  the  boric  acid  has  already  entered 
into  combination.  The  amount  of  iodine  thus  liberated  by 
the  boric  acid  is,  however,  not  large,  and  if  upon  the  addition 
of  the  iodide  and  iodate,  the  iodine  thrown  out  by  the  free 
hydrochloric  acid  present  is  immediately  bleached  by  thio- 
sulphate and  the  analysis  proceeded  with  from  this  as  the 
neutral  point,  even  in  concentrated  solutions  the  error  is 
almost  inappreciable.  If,  however,  considerable  time  inter- 
venes between  the  adding  of  the  iodide  and  iodate  and  the 
determination  of  the  neutral  point  by  thiosulphate,  as  much 
as  several  milligrams  of  boric  acid  may  have  liberated  its 
amount  of  iodine  and  is,  therefore,  not  capable  of  being 
registered  by  thiosulphate  after  the  addition  of  mannite. 
This  difficulty  was  not  met  with  in  those  experiments  in 
which  the  iodide  and  iodate  were  added  at  a  dilution  little 
greater  than  that  of  the  standard  solution  used  (7.738  grm. 
per  liter),  but  in  an  attempt  to  estimate  the  boric  acid  in 
colemanite,  where  the  solution  was  kept  as  concentrated  as 
possible,  hoping  in  this  way  to  decrease  the  time  required  for 
the  complete  liberation  of  iodine,  low  values  were  obtained; 
that  is,  a  false  starting  point  was  used. 

The  dilution  found  most  convenient  at  the  time  of  adding 
the  iodide  and  iodate  is  not  less  than  25  cm3  for  each  decigram 
of  boric  acid  (B2O3)  present  and  should  not  be  much  greater 
than  two  or  three  times  that  amount.  This  limitation  as 
regards  volume  is  equally  applicable,  whether  after  obtaining 
the  neutral  point  and  treating  with  mannite,  the  boric  acid  is 
to  be  measured  by  a  standard  solution  of  alkali  as  before 
described  or  as  here  by  the  iodine  liberated.  As  has  been 
suggested,  a  large  volume,  even  though  saturated  with  mannite, 
prolongs  the  time  of  standing  necessary  and  increases  the 
effect  of  carbon  dioxide  upon  the  iodide  and  iodate  present,  for 
carbon  dioxide,  whether  derived  from  the  atmosphere  or 
existing  dissolved  in  the  solution,  upon  standing,  slowly 
liberates  iodine.  The  amount,  however,  is  small,  and,  in  the 
time  required  for  the  completion  of  the  process,  has  never  been 


248  AN  IODOMETRIC  METHOD  FOR 

found  equivalent  to  more  than  a  single  drop  of  the  solution  of 
thiosulphate  used.  Even  if  the  material  to  be  analyzed 
contains  carbonates,  after  acidifying  in  concentrated  solution 
and  shaking  vigorously,  the  small  amount  of  uncombined 
carbon  dioxide  remaining  has  almost  an  inappreciable  effect 
upon  the  results.  The  length  of  tune  required  for  the 
liberation  of  the  theoretical  amount  of  iodine  hi  a  solution  of 
the  volume  suggested  above,  is  20  to  45  minutes,  and  at  the  end 
of  45  minutes  standing  in  a  solution  saturated  with  mannite  the 
reaction  may  be  considered  complete.  During  this  period,  how- 
ever, it  is  well  to  keep  the  solution  cool  —  at  zero  will  do  no  harm 
—  and  shake  occasionally  to  insure  thorough  mixture.  The  free 
iodine  would  tend  to  escape  upon  standing  unless  kept  in  a 
closed  flask,  but  it  is  more  convenient,  immediately  after  the 
addition  of  mannite,  to  treat  with  an  excess  of  the  standard 
solution  of  thiosulphate  —  8  or  10  cm3  more  than  the  amount 
required  to  bleach  the  iodine  liberated,  and  at  the  expiration 
of  40  to  60  minutes  titrate  back  with  ^  iodine.  The  strength 
of  the  thiosulphate  solution  found  most  convenient  is  -£",  while 
the  use  of  iodine  of  one-half  this  strength  (^£)  enables  the 
error  of  reading  to  be  correspondingly  diminished.  In  solutions 
of  the  volume  recommended  the  addition  of  starch  to  give  the 
indication  with  iodine  is  unnecessary  and  even  detrimental, 
since  a  single  drop  of  one-twentieth  normal  iodine  in  excess  is 
sufficient  to  give  a  strong  lemon  coloration,  while  in  the 
presence  of  starch  an  indefinite  dirty  red  first  appears  and 
remains  until  the  blue  is  brought  out  by  the  further  addition 
of  iodine. 

With  these  observations  in  mind,  a  series  of  experiments  was 
made  in  which  the  standard  solution  of  boric  acid  was  drawn 
into  an  Erlenmeyer  flask,  containing  a  small  amount  of  free 
hydrochloric  acid  and  made  up  to  a  definite  volume.  To 
bring  the  conditions  to  those  of  an  actual  analysis  about  1  grm. 
of  crystalline  calcium  chloride  in  solution  was  also  added. 
Potassium  iodate  (5-10  cm3  of  a  5  per  cent  solution)  and 
iodide  (3-5  cm3  of  a  40  per  cent  solution)  were  added,  and  the 
iodine  liberated  by  the  hydrochloric  acid,  barely  bleached  and 


THE  ESTIMATION  OF  BORIC  ACID. 


249 


again  brought  to  coloration  by  iodine.  Mannite  was  added  to 
saturate  the  solution,  an  excess  of  standard  thiosulphate  put 
in,  and  the  solution  set  aside  for  various  periods  of  time,  at 
the  end  of  which  the  excess  of  thiosulphate  was  titrated  by 
iodine  and  the  amount  of  unrecovered  thiosulphate  taken  as  a 
measure  of  the  boric  acid  present. 

The   thiosulphate  used  was  0.198  normal  and  the  iodine 
0.0996  normal.     The  solution  of  boric  acid  contained  7.733 

grm.  per  liter. 

TABLE  II. 


B308 

taken. 

Thio. 
taken. 

Iodine 
taken. 

Time  of 
standing. 

Volume. 

S&. 

B203 

found. 

Error. 

A 

cm3 
28.00 
27.03 
27.02 

cm3 
32.00 
32.00 
31.97 

cm8 
1.88 
4.37 
4.04 

hrs. 

0.30 
0.27 
1.00 

cm3 
28 
27 

27 

grm. 
0.2165 
0.2090 
0.2089 

gnu. 

0.2168 
0.2081 
0.2090 

gnu. 

0.0003+ 
0.0009- 
0.0001+ 

B 

27.06 
27.02 
27.04 

32.04 
32.02 
31.72 

3.88 
4.40 
3.39 

1.00 
1.00 
1.00 

50-60 
50-60 
50-60 

0.2093 
0.2089 
0.2091 

0.2101 
0.2081 
0.2096 

0.0008+ 
0.0058- 
0.0005+ 

C 

27.01 
26.05 

31.53 
31.01 

2.88 
4.01 

2.00 
3.00 

50-60 
50-60 

0.2089 
0.2014 

0.2100 
0.2025 

0.0011+ 
0.0011+ 

D 

27.00 
27.00 
26.01 
27.03 
27.05 
26.07 
27.00 

31.00 
32.00 
32.02 
31.01 
31.89 
31.02 
32.04 

2.12 
4.05 
6.20 
2.21 
3.81 
4.14 
4.30 

0.30 
0.30 
0.30 
0.48 
0.45 
0.40 
0.40 

50-60 
50-60 
50-60 
50-60 
50-60 
50-60 
60 

0.2088 
0.2088 
0.2011 
0.2090 
0.2092 
0.2016 
0.2088 

0.2089 
0.2092 
0.2018 
0.2087 
0.2093 
0.2020 
0.2086 

0.0001+ 
0.0004+ 
0.0007+ 
0.0003- 
0.0001+ 
0.0004+ 
0.0002- 

These  results  are  so  regular  that  the  method  seems  worthy  of 
high  commendation,  and  especially  since  the  standard  solutions, 
thiosulphate  and  iodine,  upon  which  the  process  depends,  are 
so  easily  prepared  and  generally  at  hand. 

The  full  method  of  procedure  recommended  is  as  follows : 
The   borate  is   dissolved  in  as   small  volume   and  as  little 


250 


AN  IODOMETRIC  METHOD  FOR 


hydrochloric  acid  as  possible,  shaking  well  to  remove  free 
carbon  dioxide  and  diluting  so  that,  at  the  time  of  adding 
potassium  iodide  and  iodate,  there  shall  be  approximately  25- 
50  cm3  of  solution  for  each  decigram  of  boric  anhydride 
present.  The  greater  part  of  the  excess  of  hydrochloric  acid 
in  the  solution  is  destroyed  by  sodium  hydroxide  and  the  use 
of  litmus  paper,  leaving  the  solution  distinctly  acid  in  reaction. 
Potassium  iodide  (3-5  cm3  of  a  40  per  cent  solution),  and 
iodate  (5-10  cm3  of  a  5  per  cent  solution)  are  added  in  excess 
of  that  required  to  liberate  iodine  in  an  amount  corresponding 
to  the  hydrochloric  acid  and  the  boric  acid  present.  The 
iodine  liberated  by  the  free  hydrochloric  acid  is  bleached  by  a 
small  amount  of  a  strong  solution  of  thiosulphate,  and  after 
agitating  to  insure  thorough  mixture,  iodine  is  added  to  faint 
coloration.  Sufficient  mannite  is  now  used  to  saturate  the 
solution  —  about  10-15  grm.  for  a  volume  of  50  cm3  —  and 
sodium  thiosulphate  added  in  standard  solution  8-10  cm8  in 
excess  of  that  required  to  bleach  the  iodine  immediately  thrown 
out  by  the  mannite.  The  solution  is  again  brought  to 
saturation,  if  necessary,  by  mannite  and  after  standing  in  a 
cool  place  for  40-60  minutes,  titrated  with  decinormal  iodine 
to  determine  the  excess  of  thiosulphate  present.  In  the 
manner  described,  specimens  of  crude  calcium  borate  and 
crystals  of  colemanite  were  analyzed  with  the  results  given 
below. 

TABLE  III. 
CALCIUM  BORATE. 


Mineral. 

Thio. 
taken. 

Iodine 
taken. 

Time 
stand- 
ing. 

Volume 
of  solu- 
tions. 

B20. 

found. 

Per  cent. 

grm. 
0.4015 
0.4010 

cms 
35.05 
35.34 

cm8 
4.75 
5.23 

hrs. 
1.00 
2.00 

cm8 

40 
45 

grm. 
0.2280 
0.2283 

56.92 
56.94 

COLEMANITE. 

0.4002 
0.2513 
0.4007 

32.00 
32.01 
33.03 

5.50 
7.36 

7.72 

1.30 
1.00 
0.50 

50 
40 
65 

0.2043 
0.1279 
0.2036 

51.04 
50.91 
50.81 

THE  ESTIMATION  OF  BORIC  ACID.  251 

The  solution  of  thiosulphate  used  was  0.19939  and  the 
iodine  0.0996  normal. 

These  results  show  little  variation  and  in  the  case  of 
colemanite  correspond  closely  to  the  theory  50.97  per  cent. 
The  process  is  convenient,  generally  applicable,  and  accurate 
within  the  ordinary  limits  of  analysis. 


XXXI 

THE  DOUBLE  AMMONIUM  PHOSPHATES   OP 

BERYLLIUM,  ZINC,  AND  CADMIUM  IN 

ANALYSIS. 

BY  MARTHA  AUSTIN.* 

IT  has  been  shown  f  that  the  composition  of  the  phosphate 
of  manganese  thrown  down  by  microcosmic  salt  from  the 
solution  of  a  pure  manganous  salt  contains  more  manganese 
than  belongs  to  the  ideal  ammonium  manganese  phosphate 
NH4MnPO4;  and,  further,  that  by  acting  with  ammonium 
chloride  in  proper  proportion  the  phosphate  of  manganese 
thrown  down  by  microcosmic  salt  may  be  completely  converted 
to  the  ideal  ammonium  manganese  phosphate.  Ammonium 
chloride,  likewise,  in  the  case  of  magnesium  phosphate  f  tends 
to  cause  the  replacement  of  the  metal  by  ammonia.  Indeed, 
the  replacement  here  is  readily  carried  so  far  beyond  the  point 
corresponding  to  the  normal  ammonium  magnesium  phosphate, 
NH4MgPO4,  that  the  tendency  to  form  a  salt  richer  in 
ammonia  and  poorer  in  magnesium  —  perhaps  something  like 
Mg(NH4)4(PO4)2 —  must  be  recognized. 

These  facts  suggested  an  investigation  into  the  constitution 
of  certain  other  ammonium  phosphates  with  reference  to  their 
utility  in  analytical  processes.  Of  the  elements  of  Mendele'eff's 
second  group,  beryllium,  magnesium,  zinc,  cadmium,  and  mer- 
cury are  capable  of  yielding  double  ammonium  phosphates, 
while  no  such  compounds  of  calcium,  strontium  and  barium 
have  been  described.  The  solubility  in  ammonia  of  the  double 
ammonium  phosphates  of  the  elements  of  the  former  category 

*  From  Am.  Jour.  Sci.,  viii,  206. 

t  Am.  Jour.  Sci.,  vol.  yi,  233.    This  volume,  p.  121. 

}  Am.  Jour.  Sci.,  vol.  vii,  187.    This  volume,  p.  190. 


DOUBLE  AMMONIUM  PHOSPHATES  IN  ANALYSIS.     253 

appears  to  increase  as  the  elements  of  which  they  are 
compounds  are  removed  in  the  series  from  the  beryllium, 
and,  while  the  same  is  true  of  the  simple  phosphates  of 
members  of  the  latter  category,  the  extent  of  such  solvent 
action  is  slight  comparatively.  According  to  the  work  recorded 
in  the  literature,  calcium,  barium,  and  strontium  form  individu- 
ally a  neutral  tribasic  phosphate  or  acid  phosphates  of  greater 
or  less  degree  of  acidity  according  to  the  conditions  of 
precipitation.  In  my  experience  where  salts  of  these  elements 
were  precipitated  either  with  ammonium  phosphate  or 
microcosmic  salt  hi  presence  of  varying  amounts  of  ammonium 
chloride,  or  ammonia,  or  both,  only  the  recognized  phosphates 
were  obtained.  The  effect  of  ammonium  salts  in  presence  of 
ammonia  seemed  to  promote  the  formation  of  the  tribasic  salt 
in  the  case  of  calcium  and  strontium;  barium  tends  to  form 
the  barium  acid  phosphate  almost  exclusively  even  in  the 
presence  of  ammonium  salts  and  free  ammonia.  No  double 
ammonium  phosphate  of  either  calcium,  strontium,  or  barium 
was  produced  under  any  condition.  As  is  well  known, 
mercury  does  form  an  ammonium  mercuiy  phosphate,  but  the 
salt  is  soluble  to  so  great  a  degree  in  ammonia,  ammonium 
chloride,  and  even  in  the  precipitant  itself,  that  nothing  of 
any  value  for  analytical  work  seemed  likely  to  come  from  its 
study. 

The  Ammonium  Beryllium  Phosphate. 

The  ammonium  beryllium  phosphate  has  beeen  described 
by  Roessler  *  as  a  crystalline  salt  produced  by  boiling  some 
time  in  ammoniacal  solution  the  phosphate  precipitated  by 
ammonium  phosphate,  though  the  best  results  of  this  treatment 
failed  to  yield  the  ideal  constitution  of  this  salt,  NH4BePO4. 
This  same  precipitate  cannot  be  obtained,  Roessler  further 
states,  by  using  a  sodium  salt  as  the  precipitant.  In  order  to 
follow  out  this  work  of  Roessler,  a  solution  of  berryllium 
chloride  for  use  was  prepared  as  follows :  The  pure  beryllium 
chloride  of  commerce  was  dissolved  in  as  little  water  as 
*  Fresenius,  Zeitschr.  anal.  Chem.,  xvii,  148. 


254    DOUBLE  AMMONIUM  PHOSPHATES  OF  BERYLLIUM, 

possible  and  treated  for  the  precipitation  of  aluminum  by 
ethereal  hydrochloric  acid.*  After  filtering  and  evaporating 
from  the  filtrate  the  ether  and  a  part  of  the  hydrochloric  acid, 
the  beryllium  was  precipitated  with  ammonia,  filtered  to 
remove  any  members  of  the  magnesium  group,  and  washed 
free  from  ammonium  chloride.  The  larger  part  of  the 
precipitate  was  dissolved  in  hydrochloric  acid  in  slight  excess, 
and  boiled  with  the  reserved  portion.  After  filtering,  the 
solution  was  diluted  to  definite  volume  and  standardized  by 
precipitating  measured  portions  of  the  solution  with  ammonia, 
filtering  on  asbestos  under  pressure  in  a  perforated  platinum 
crucible,  igniting  the  residue  and  weighing  as  the  oxide.  The 
results  recorded  in  section  A  of  the  following  table  were 
obtained  by  precipitating  definite  volumes  of  the  pure  solution 
of  beryllium  chloride  with  ammonium  phosphate  in  a  platinum 
dish,  dissolving  the  precipitate  in  hydrochloric  acid  in  faint 
excess,  and  while  hot  precipitating  slowly  with  dilute  ammonia, 
boiling  (while  the  solution  was  kept  distinctly  ammoniacal) 
until  the  flocky  precipitate  was  entirely  converted  to  a  fine, 
powdery,  semi-crystalline,  rapidly  subsiding  mass.  A  quarter 
to  a  half-hour  is  necessary  under  the  most  favorable  conditions 
to  cause  this  conversion.  After  cooling,  the  precipitate  was 
filtered  off  on  asbestos  under  pressure  in  a  perforated  platinum 
crucible,  washed  carefully  with  distilled  water,  dried,  ignited 
and  weighed.  The  filtrate  was  tested  for  beryllium  by  boiling 
with  ammonia.  None  was  found  in  these  cases,  nor  in  any  of 
the  following  work.  Faint  traces  of  chloride  were  found  in 
the  residues  after  ignition  after  dissolving  in  nitric  acid  and 
testing  with  silver  nitrate. 

The  results  are  in  every  case  in  excess  of  the  theory  for  the 
pyrophosphate  derived  by  ignition  of  the  ammonium  beryl- 
lium phosphate,  possibly  because  the  ammonium  chloride 
present  may  have  a  tendency  to  form  a  salt  too  rich  in  ammo- 
nium (as  was  shown  to  be  the  case  with  the  magnesium  salt), 
consequently  giving  too  much  phosphoric  acid  in  the  ignited 
residue ;  or,  because  of  inclusion  of  the  chloride  and  phosphoric 

*  Am.  Jour.  Sci.,  ir,  111.    This  volume,  p.  111. 


ZINC,  AND  CADMIUM  IN  ANALYSIS.  255 

acid.  It  might  reasonably  be  expected  that  some  phosphoric 
acid  may  be  held,  since  a  trace  of  chloride  was  found.  Either 
or  both  of  these  substances  may  have  been  held  mechanically, 
or  in  combination. 

It  was  found  that  on  boiling  for  some  tune  the  solution  of 
beryllium  chloride  with  microcosmic  salt — (6)  section  B  of 
the  table  —  and  precipitating  in  the  same  manner  as  when 
ammonium  phosphate  was  used  the  same  sort  of  powdery 
mass  remained  as  was  obtained  by  the  ammonium  phosphate. 
The  residue  being  tested  for  sodium  according  to  the  method 
brought  out  by  Kreider  and  Brecken^idge,*  showed  sodium 
present  to  the  amount  of  0.0062  grm.  reckoned  as  sodium 
phosphate.  It  may  be  reasonably  supposed  that  the  presence 
of  the  sodium  was  due  to  one  of  two  causes,  —  inclusion  of 
the  soluble  phosphate,  or  a  tendency  on  the  part  of  the  beryl- 
lium to  form  an  ammoniumf  sodium  beryllium  phosphate  or  a 
sodium  J  beryllium  phosphate,  both  of  which  are  known  to 
exist.  Long  boiling  of  the  precipitates  is  tedious,  and,  unless 
great  care  is  taken,  may  involve  small  losses  of  material ;  hence 
if  the  same  results  could  be  obtained  with  less  boiling  such 
treatment  would  be  decidedly  advantageous.  The  results  in 
section  C  of  the  table  were  obtained  by  adding  microcosmic 
salt  to  the  hot  solutions  of  the  chloride,  boiling  five  minutes, 
cooling,  filtering  off  on  an  ashless  filter  —  because  of  the  flocky 
condition  of  the  precipitate  —  treating  as  usual  before  igniting 
the  residue  in  a  platinum  crucible.  The  results  compare  well 
with  those  obtained  by  long  boiling  of  the  precipitated  beryl- 
lium —  although  all  are  in  excess  of  the  theory.  That  ammo- 
nium chloride  here,  as  in  cases  above,  has  a  marked  effect  in 
changing  the  constitution  of  the  phosphate  precipitated  by 
microcosmic  salt  is  not  readily  seen.  It  is  obvious  that  the 
presence  of  an  excess  of  the  soluble  phosphate  is  essential  to 
precipitate  the  beryllium  as  the  double  ammonium  phosphate 
from  the  results  recorded  in  section  D  of  the  table,  where, 

*  Am.  Jour.  Sci.,  ii,  263.    Volume  I,  p.  401. 

t  Persoz,  Ann.  Chem.  (Liebig),  Ixv,  174 ;  Atterberg,  Bulletin  Soc.  Chim., 
xxiv,  358. 

J  Scheffer,  Ann.  Chem.  (Liebig),  cix,  144. 


256    DOUBLE  AMMONIUM  PHOSPHATES  OF  BERYLLIUM, 


after  the  precipitate  of  beryllium  phosphate  had  subsided  and 
the  supernatant  liquid  had  been  poured  off,  the  precipitate  dis- 
solved in  hydrochloric  acid  was  brought  down  again  at  the 
boiling  temperature  with  ammonia  either  alone  or  in  presence 
of  ammonium  chloride.  The  results  obtained  show  that  the 
salt  approaches  the  constitution  of  the  tribasic  phosphate, 
when  it  is  precipitated  in  presence  of  a  faint  excess  of  phos- 
phoric acid,  even  though  ammonium  chloride  in  large  amount 
be  present. 

TABLE  I. 


Exp. 

Be2P2O7  corresponding 
to  BeCl,. 

Be3P,O8  corresponding 
to  BeCl,. 

(NH4)3P04. 

NB^Cl. 

Taken. 

Found. 

Error. 

Taken. 

Found. 

Error. 

A. 

(1) 
(2) 
(3) 
(4) 
(5) 

grm. 
0.3578 
0.3578 
0.3578 
0.3578 
0.3578 

grin. 

0.3613 
0.3808 
0.3707 
0.3640 

0.3680 

grin. 

0.0035+ 
0.0230+ 
0.0129+ 
0.0062+ 
0.0102+ 

grm 

grm. 

grm. 

grm. 
2 
2 
2 
2 
2 

grm 

'  'so 

B. 

HNaNH4PO4 
.4H26. 

(6) 

0.3578 

0.3697  |  0.0119+  |    ... 

...    |      ... 

C. 

(7) 
(8) 
(9) 
(10) 

0.3578 
0.3578 
0.3578 
0.3578 

0.3618 
0.3680 
0.3729 
0.3631 

0.0040+ 
0.0102+ 
0.0151+ 
0.0053+ 

1.2 
1.2 
1.2 
1.2 

'  *10 
60 

D. 

(11) 
(12) 
(13) 
(14) 

0.2700 
0.2700 
0.2700 
0.2700 

0.2589 
0.2989 
0.2936 
0.2507 

0.0111- 
0.0289+ 
0.0236+ 
0.0193- 

0.5 
0.5 
0.5 
0.5 

—10 
5-60 
—60 

From  the  work  described  it  is  clear  that  the  ammonium 
beryllium  phosphate  is  not  obtained  in  ideal  condition  by  pre- 
cipitating a  solution  of  the  chloride  with  ammonium  phosphate. 
Roessler's  own  results  were  likewise  only  approximately  cor- 
rect, as  he  states.  It  is  also  plain  that  hydrogen  sodium 
ammonium  phosphate  precipitates  the  ammonium  beryllium 


ZINC,  AND   CADMIUM  IN  ANALYSIS.  257 

phosphate  in  a  condition  as  nearly  ideal  as  does  the  ammonium 
phosphate,  while  the  effect  of  the  ammonium  chloride  in  either 
case  is  not  marked  in  producing  a  phosphate  containing 
ammonia.  Of  most  importance  in  obtaining  the  ammonium 
salt  is  an  excess  of  the  soluble  phosphate,  for  when  the  amount 
of  the  precipitant  is  reduced  to  a  little  more  than  the  theo- 
retical amount  the  condition  of  the  phosphate  coincides  almost 
exactly  with  the  theory  for  the  tribasic  phosphate,  even  though 
a  large  excess  of  ammonium  chloride  be  present.  When  there 
is  an  abundance  of  the  precipitant  the  results  are  all  in  excess 
of  the  theory,  which  may  be  accounted  for  on  the  supposition 
that  foreign  material  is  included  —  the  chloride  of  ammonia  and 
the  soluble  phosphate  —  to  a  greater  or  less  extent  by  the  pre- 
cipitate. The  formation  of  a  phosphate  of  beryllium  contain- 
ing too  much  ammonia  and  phosphoric  acid,  or,  in  case  of  the 
precipitations  by  microcosmic  salt,  sodium  by  the  formation 
of  a  sodium  ammonium  beryllium  phosphate  and  sodium  beryl- 
lium phosphate  (known  salts),  is  not  definitely  proved. 

The  Ammonium  Zinc  Phosphate. 

Debray,*  Bette  f  and  Heintz  J  separately  found  that  am- 
monium zinc  phosphate  is  formed  by  boiling  a  solution  of 
zinc  sulphate  with  ammonium  phosphate.  This  salt  was 
investigated  later  by  A.  Guyard  (Hugo  Tamm),§  who  found 
that  if  to  a  solution  of  a  zinc  salt  of  an  organic  or  a  mineral 
acid  supersaturated  with  ammonia  until  all  the  zinc  oxide 
is  dissolved  and  made  faintly  acid  with  hydrochloric  acid, 
sodium  phosphate  be  added,  a  flocky  precipitate  resulted, 
which  on  being  kept  near  the  boiling  point  for  some  seconds 
was  converted  to  crystalline  zinc  ammonium  phosphate,  which 
filtered  readily  and  was  washed  free  from  impurities  with  the 
greatest  facility.  He  found  that  all  the  zinc  in  solution  was 
thrown  down  as  the  ammonium  zinc  phosphate,  which  on 
ignition  yielded  the  zinc  pyrophosphate.  With  care  in 
handling  this  process  to  avoid  an  excess  of  the  precipitant, 

*  Compt.  rend.,  lix,  40.  t  Ann.  Chem.  (Liebig),  xv,  129. 

J  Ann.  Chem.  (Liebig),  cxliii,  156.      §  Chem.  News,  xxir,  148. 
VOL.  ix.  — 17 


258    DOUBLE  AMMONIUM  PHOSPHATES  OF  BERYLLIUM, 

and  the  presence  of  sodium  and  potassium  salts  (on  account 
of  the  danger  of  occlusion)  the  precipitation  of  the  ammo- 
nium zinc  phosphate,  ignition,  and  weighing  as  the  pyro- 
phosphate  made,  Guyard  believed,  an  ideal  process  for  the 
estimation  of  zinc.  Although  there  was  slight  solubility  of 
the  salt,  it  made  an  insignificant  loss  when  the  process  was 
handled  properly.  Acids  present,  or  certain  alkalies  to  any 
great  extent,  increased  the  solubility  of  the  salt  so  much 
that  the  loss  became  appreciable.  Another  source  of  error 
was  to  Guyard's  mind  loss  of  zinc  during  the  ignition 
of  the  zinc  ammonium  phosphate  with  the  paper  on  which 
the  precipitate  had  been  collected.  Garrigues*  found,  in 
estimating  zinc  in  a  practical  way,  that  this  process  advocated 
by  Guyard  gives  in  solutions  of  zinc  free  from  salts  of  all 
metals,  even  alkaline  salts  —  solutions  that  from  previous 
steps  in  analysis,  however,  must  have  contained  ammonium 
chloride  in  large  amount  —  as  satisfactory  results  as  Guyard 
claimed  for  it.  Garrigues'  method  of  procedure  was  to  add 
acid  diammonium  phosphate  to  a  warm  solution  of  zinc 
exactly  neutralized  with  either  hydrochloric  acid  or  ammonia, 
so  that  the  weights  of  zinc  ammonium  phosphate  and  that 
of  the  diammonium  phosphate  added  should  be  as  one  to 
five  respectively,  to  heat  until  the  flocky  precipitate  becomes 
crystalline  and  subsides,  filtering  off  on  asbestos,  drying  at 
100°  C.  and  weighing  preferably,  although  the  residue  may 
be  ignited  without  loss,  since  the  filtration  is  made  on  asbestos 
in  a  perforated  crucible.  Langmuir  f  modifies  the  method  by 
destroying  with  dilute  acetic  acid  any  free  ammonia  that  may 
be  left  in  the  solution  after  boiling. 

In  the  work  that  follows,  in  which  an  attempt  was  made  to 
show  what  precipitate  is  formed  from  a  solution  of  zinc  by  the 
action  of  a  soluble  phosphate,  also  what  effect  ammonium 
chloride  has  upon  the  precipitate,  a  solution  of  zinc  chloride 
prepared  as  detailed  below  was  employed.  The  pure  zinc 
chloride  of  commerce  was  treated  with  zinc  carbonate,  filtered 
and  precipitated  with  ammonium  sulphide.  This  precipitate 

*  Jour.  Am.  Chem.  Soc.,  xix,  936.          t  Jour.  Am.  Chem.  Soc.,  xxi,  115. 


ZINC,  AND  CADMIUM  IN  ANALYSIS. 


259 


was  boiled  in  a  slight  excess  of  hydrochloric  acid  until  all  the 
hydrogen  sulphide  was  removed,  and  then  was  precipitated 
with  sodium  carbonate.  After  washing  carefully  until  all 
the  chloride  was  removed,  the  greater  part  of  the  carbonate 
was  dissolved  in  sulphuric  acid  in  slight  excess,  boiled  with 
the  remaining  portion  of  the  carbonate  and  filtered.  This 
solution  diluted  to  definite  volume  was  standardized  as  sul- 
phate by  evaporating  the  solution  to  dryness  in  a  platinum 
crucible  and  heating  the  residue.*  The  heating  is  carried 
on  safely  by  so  placing  the  platinum  crucible  in  a  radiator 
(consisting  of  a  crucible  and  a  triangle)  that  the  bottom  of 
the  platinum  crucible  was  held  about  one  centimeter  above 
the  bottom  of  the  outside  crucible.  Constant  weights  were 
obtained  in  successive  treatment  with  a  few  drops  of  sulphuric 
acid  and  heating  over  the  radiator.  The  results  obtained  in 
this  manner  were  a  trifle  higher,  though  in  fair  agreement 
(when  the  nature  of  the  carbonate  process  is  taken  into 
consideration)  with  determinations  of  the  zinc  in  the  solu- 
tions as  oxide  after  precipitating  with  sodium  carbonate  with 
the  usual  precautions,  filtering  off  on  asbestos  under  pressure 
in  a  perforated  platinum  crucible,  washing  with  distilled 
water,  drying,  and  igniting.  Results  are  given  in  Table  II 
showing  the  amount  of  zinc  sulphate  found  in  five  different 
portions  each  of  forty  cubic  centimeters  of  the  solution 
of  zinc  sulphate,  and,  for  comparison,  the  results  of  de- 
terminations as  zinc  oxide  by  the  carbonate  processes  are 

included. 

TABLE  II. 


ZnSO4  found  in 
40  cm»  of 
solution. 

Mean  value  of  ZnO 
corresponding  to 
ZnSO4  in  40  cms 
of  solution. 

ZnO  found  in  40  cm8 
of  solution  by 
precipitation  as 
the  carbonate. 

grm. 

grim* 

gnu. 

0.5386  ' 

0.2712 

0.2691 

0.5385 

0.2685 

0.5387 

0.2711 

0.5387 

t 

0.5390  j 

Rose-Finkener,  Analytische  Chemie,  6te  Auflage,  vol.  ii,  117. 


260    DOUBLE  AMMONIUM  PHOSPHATES  OF  BERYLLIUM, 

Definite  portions  of  the  solution  of  zinc  sulphate  were 
carefully  drawn  from  a  burette  into  a  platinum  dish,  heated  and 
treated  with  ammonium  phosphate  until  the  solution  turned 
red  litmus  paper  blue.  The  whole  was  heated  until  the  flocky 
precipitate  became  crystalline  and  fell  to  the  bottom  of  the 
dish.  The  solution  after  standing  as  recorded  in  section  A  of 
the  table  was  filtered  off  on  asbestos  under  pressure  in  a 
perforated  platinum  crucible,  and  the  precipitate  was  washed 
with  distilled  water,  dried,  ignited  and  weighed.  The  filtrate 
in  each  case,  as  in  all  following  cases,  was  tested  for  zinc  with 
sulphuretted  hydrogen.  The  results  recorded  in  section  B  of 
the  table  were  obtained  in  the  same  manner  as  those  of  section 
A,  with  microcosmic  salt  substituted  for  the  ammonium  salt  as 
the  precipitant.  The  results  are  below  the  theory  for  the 
pyrophosphate,  but  no  appreciable  amount  of  zinc  appeared  in 
the  filtrates.  Neither  ammonium  phosphate  nor  ammonium 
sodium  phosphate  seems  to  precipitate  the  ideal  ammonium 
zinc  phosphate  under  these  conditions;  and  the  time  of 
standing  appears  to  be  without  effect. 

The  results  recorded  in  section  C  were  obtained  by  precipi- 
tating the  warm  solution  of  the  zinc  in  presence  of  large 
amounts  of  ammonium  chloride  by  adding  microcosmic  salt 
until  the  solution  was  alkaline  to  litmus.  From  these  results 
it  seems  that  the  presence  of  ammonium  chloride  is  essential 
for  the  conversion  of  the  zinc  phosphate  precipitated  by 
hydrogen  sodium  ammonium  phosphate  to  the  ammonium  zinc 
salt.  As  a  matter  of  fact  the  solutions  employed  by  Guyard 
and  those  in  which  estimations  are  made  by  practical  workers 
do  contain  ammonium  chloride  formed  in  previous  steps  of  the 
analysis.  The  proportion  of  zinc  to  phosphate  suggested  by 
Garrigues  —  1  :  5  —  is  the  amount  of  soluble  phosphate  neces- 
sary to  turn  red  litmus  blue  after  the  zinc  is  precipitated.  In 
order  to  find  out  whether  the  presence  of  so  large  an  amount 
of  the  soluble  phosphate  is  necessary  in  presence  of  ammonium 
chloride,  the  solution  of  zinc  sulphate  was  precipitated  in 
presence  of  the  necessary  amount  of  ammonium  chloride  by 
the  microcosmic  salt  in  small  excess  above  the  equivalent  of 


ZINC,  AND  CADMIUM  IN  ANALYSIS. 


261 


the  zinc  salt,  and  the  solution  was  made  just  ammoniacal 
to  litmus  with  a  few  drops  of  dilute  ammonia  both  before 
and  after  heating  to  convert  the  precipitate  to  crystalline 
condition.  Experiment  (15)  shows  that  precipitation  is  not 
complete  under  these  conditions.  The  zinc  left  in  the  solution 
was  precipitated  at  once  as  sulphide,  and  estimated  as  the 
oxide,  after  dissolving  in  hydrochloric  acid  and  precipitating 

TABLE  III. 


Exp. 

Zn2P2O7 
corre- 
sponding 
to 
ZnS04. 
Taken. 

Pound. 

Error. 

> 

Error 
in  terms 
of  Zinc. 

Zn2P207 
corre- 
sponding 
to  Zn 
left  in  the 
filtrate. 

(NH4)^04. 

NH4CL 

Time  of 
stand- 
ing* 

A. 

$ 

(3) 

grm. 
0.6355 
0.6355 
0.6355 

grin. 

0.6206 
0.6254 
0.6300 

grin. 

0.0149- 
0.0101- 
0.0055- 

grm. 
0.0060- 
0.0040- 
0.0022- 

grm. 
Trace. 
Trace. 
Trace. 

grm. 

3.13 
3.13 
3.13 

grm. 

hrs. 

iJ* 

16 

B. 

(4) 
(5) 

0.6355 
0.6355 

0.6271 
0.6256 

0.0084- 
0.0099- 

0.0034- 
0.0040- 

Trace. 
None. 

HNaNH4P04 
.  4HaO. 

0.5 
0.5 

1 

20 

4.47 
4.47 

C. 

(6) 

| 

(9). 
(10) 

(11) 
(12) 
(13) 
(14) 

0.6355 
0.6355 
0.6355 
0.6355 
0.6355 
0.6355 
0.6355 
0.6355 
0.6367 

0.6285 
0.6304 
0.6295 
0.6335 
0.6381 
0.6379 
0.6386 
0.6393 
0.6355 

0.0070- 
0.0051- 
0.0060- 
0.0020- 
0.0026+ 
0.0024+ 
0.0031+ 
0.0038+ 
0.0012+ 

0.0028- 
0.0020- 
0.0024- 
0.0008- 
0.0010+ 
0.0009+ 
0.0012+ 
0.0014+ 
0.0005+ 

None. 
None. 
None. 
None. 
None. 
None. 
None. 
None. 
None. 

4.47 
4.47 
4.47 
4.47 
4.47 
4.47 
4.47 
4.47 
4.47 

10 
10 
10 
10 

20 
20 
20 
20 
30 

16 

,> 

j 

D. 

(15) 
(16) 

0.6355 
0.6355 

0.6172 
10.6227 
II  0.0040 

0.0183- 
0.0098- 

0.0072- 
0.0039- 

0.0108 
None. 

0.894 
10.894 
II  3.576 

20 
20 

3 
* 

E. 

(17) 
(18) 
(19) 

0.6355 
0.6355 
0.6355 

0.6270 
0.6125 
0.6303 

0.0085- 
0.0230- 
0.0052- 

0.0034- 
0.0093- 
0.0021- 

None. 
0.0148 
0.0020 

4.47 
4.47 
4.47 

10 

3 
18 
18 

262    DOUBLE  AMMONIUM  PHOSPHATES  OF  BERYLLIUM, 

with  sodium  carbonate.  In  (16)  of  the  table  the  first  nitrate 
was  treated  with  an  excess  of  microcosmic  salt,  and  boiled. 
Another  portion  of  the  ammonium  zinc  phosphate  was 
precipitated,  and  was  filtered  off  and  estimated.  No  zinc 
was  found  by  sulphuretted  hydrogen  in  the  second  filtrate. 

From  the  results  it  seems  obvious  also  that  an  excess  of  the 
soluble  phosphate  is  necessary  to  complete  the  precipitation  of 
the  zinc  as  the  ammonium  zinc  phosphate  instead  of  partly 
ammonium  zinc  phosphate  and  partly  tribasic  phosphate. 

In  section  E  of  the  table  are  recorded  results  where  the 
precipitation  was  made  in  presence  of  an  excess  of  the  precipi- 
tant either  alone  or  in  presence  of  ammonium  chloride,  the 
solution  being  made  faintly  acid  to  litmus  with  acetic  acid, 
according  to  the  manner  in  which  Langmuir  recommends  to 
conduct  the  precipitation.  All  the  results  by  the  method  are 
low.  The  condition  of  the  ammonium  zinc  phosphate  most 
nearly  approximating  to  the  ideal  is  obtained  as  shown  in  ( 9) 
to  (14)  by  precipitating  in  presence  of  ammonium  chloride  in 
large  amount.  Microcosmic  salt  is  added  until  the  solution 
containing  the  ammonium  salt  is  alkaline  and  the  whole  is 
heated  until  the  mass  subsides  in  crystalline  condition.  The 
amount  of  ammonium  chloride  should  be  twenty  grams  if  the 
filtration  is  to  be  made  as  soon  as  the  solution  cools.  One-half 
the  amount  will  do  if  the  liquid  stands  a  number  of  hours. 
Larger  amounts  tend  to  give  a  salt  too  rich  in  ammonia.  The 
time  of  standing  seems  to  be  a  less  important  factor  than  either 
the  excess  of  microcosmic  salt  or  ammonium  chloride. 

The  Ammonium  Cadmium  Phosphate. 

According  to  S.  Drewsen*  the  cadmium  ammonium  phos- 
phate is  precipitated  by  allowing  a  solution  of  cadmium 
sulphate  to  stand  twenty-four  hours  with  ammonium  phosphate. 
It  is  very  soluble  both  in  acids  and  alkalies.  No  further 
preparation  of  this  seems  to  have  been  recorded.  For  the 
work  on  this  salt  to  be  given  below,  done  with  reference  to 
the  constitution  of  the  salt  precipitated  by  hydrogen  sodium 

*  Gmelin-Kraut,  6te  Auflage,  iii,  74. 


ZINC,  AND  CADMIUM  IN  ANALYSIS.  263 

ammonium  phosphate,  the  effect  of  ammonium  chloride  in  the 
precipitation,  and  the  value  of  the  salt  for  quantitative  work, 
the  solution  of  cadmium  chloride  employed  was  prepared  as 
follows :  A  solution  of  cadmium  sulphate  acidulated  with 
hydrochloric  acid  was  precipitated  with  sulphuretted  hydrogen, 
filtered  and  washed,  and  the  precipitated  sulphide  was  dissolved 
in  hydrochloric  acid  and  filtered  from  possible  traces  of  copper 
and  lead.  The  solution  of  the  sulphide  in  hydrochloric  acid 
was  boiled  until  all  the  sulphuretted  hydrogen  was  expelled, 
and  filtered  on  asbestos  in  a  perforated  crucible  of  platinum 
under  pressure.  The  cadmium  in  the  filtrate  precipitated  with 
ammonium  carbonate  in  excess  was  washed  free  from  chloride, 
dissolved  in  hydrochloric  acid  and  diluted  to  definite  volume. 
It  was  standardized  as  oxide*  after  precipitating  with  sodium 
carbonate  with  the  necessary  precautions. 

The  standard  solution  of  cadmium  chloride  was  drawn 
carefully  from  a  burette  into  a  platinum  dish,  and,  while  hot, 
was  precipitated  by  adding  hydrogen  sodium  ammonium 
phosphate  until  the  solution  was  alkaline  to  litmus.  After 
heating  until  the  solution  became  crystalline,  the  whole  stood 
three  hours  in  case  of  (1)  of  the  table  and  sixteen  hours  in 
case  of  (2)  and  (3),  before  filtering.  In  experiments  (4)  to 
(12),  inclusive,  recorded  in  the  table,  precipitation  was  made 
in  the  same  manner  as  in  (1)  to  (3)  in  presence  of  varying 
amounts  of  ammonium  chloride,  and  the  precipitates  were 
filtered  after  standing  as  stated  below  in  the  table.  It  is  clear 
from  the  results  that  the  cadmium  separates  out  completely  on 
long  standing  only.  Moreover,  the  ideal  condition  of  the 
ammonium  cadmium,  phosphate  is  obtained  only  when  an 
abundance  of  ammonium  chloride  is  present;  but  large 
amounts  of  ammonium  chloride  dissolve  this  salt.  In  (14), 
where  ammonia  was  added  after  precipitation  was  complete, 
the  salt  dissolved  somewhat ;  also  in  (15),  where  the  solution 
was  left  faintly  acid  with  acetic  acid,  a  large  part  of  the  salt 
was  dissolved.  These  weights  of  cadmium  dissolved  in  the 
filtrate  were  obtained  by  treating  the  filtrates  with  sulphuretted 
*  Browning,  Am.  Jour.  Sci.,  xlvi,  280.  Volume  I,  p.  226. 


264    DOUBLE  AMMONIUM  PHOSPHATES  OF  BERYLLIUM. 


TABLE  IV. 


Cd2P,07 

CdsP207 
corre- 

Time 

Exp. 

corre- 
sponding 
to  CdCl,. 

Found. 

Error. 

Errorin 
terms  of 
Cadmium. 

sponding 
toCd 
found  in 

HNaNEUPO* 
.iHsO. 

NH4C1. 

of 
stand 
ing. 

Taken. 

the  nitrate. 

griu. 

grm. 

grm. 

grm. 

grm. 

grm. 

grm. 

hra. 

(1) 

0.6972 

0.6201 

0.0771- 

0.0434- 

0.0059 

4.5 

t 

3 

(2) 

0.6972 

0.6135 

0.0837- 

0.0471- 

None. 

4.5 

.  . 

16 

(3) 

0.6972 

0.6134 

0.0838- 

0.0471- 

None. 

4.5 

.  . 

16 

(4) 

0.6972 

0.6792 

0.0180- 

0.0101- 

Trace. 

4.6 

1 

16 

(5) 

0.6972 

0.6831 

0.0141- 

0.0079— 

0.0113 

4.5 

10 

2 

(6) 

0.6972 

0.6976 

0.0004+ 

0.0002-f 

Trace. 

4.5 

10 

16 

(7) 

0.6972 

0.6969 

0.0003- 

0.0002- 

Trace. 

4.5 

10 

18 

(8) 

0.6972 

0.6962 

0.0010- 

0.0006- 

Trace. 

4.5 

10 

16 

9 

0.6972 

0.6891 

0.0081- 

0.0045- 

0.0191 

4.5 

20 

16 

(10) 

0.6972 

0.6972 

0.0000 

0.0000 

Trace. 

4.5 

20 

16 

(11) 

0.6972 

0.6942 

0.0030- 

0.0016- 

Trace. 

4.5 

20 

16 

(12) 

0.6972 

0.6737 

0.0235- 

0.0132- 

0.0304 

4.5 

30 

16 

(13) 

0.6972 

0.5655 

0.1317- 

0.0741- 

0.1378 

4.5 

30 

16 

(14) 

0.6972 

0.6922 

0.0050- 

0.0023- 

0.0088 

4.5 

10 

16 

(15) 

0.6972 

0.3209 

0.3763- 

0.2117- 

0.2449 

4.5 

•  • 

16 

hydrogen,  dissolving  the  sulphide  in  nitric  acid,  and  weighing 
as  oxide  after  precipitating  with  sodium  carbonate. 

The  ammonium  cadmium  phosphate  is  obtained  in  ideal 
condition  by  precipitating  with  microcosmic  salt  in  presence  of 
10  grm.  ammonium  chloride  in  a  total  volume  of  100  cm3  to 
150  cm3  —  shown  in  (6),  (7),  and  (8)  — filtering  after  standing 
some  time.  On  drying  and  igniting  the  pyrophosphate  is  left. 
Very  large  amounts  of  ammonium  chloride  —  30  grm.  — 
dissolve  the  salt,  and  seem  to  tend  to  cause  the  formation  of  a 
phosphate  too  rich  in  ammonia.  Either  acid  or  ammonia  in 
small  amount  dissolves  the  salt,  as  is  shown  in  (14)  and  (15). 

The  results  of  this  investigation  as  to  the  analytical 
application  of  the  double  ammonium  phosphates  of  beryllium, 
zinc,  and  cadmium  may  be  summarized  briefly  as  follows :  It 
is  impossible  to  estimate  beryllium  with  accuracy  as  the 
pyrophosphate  obtained  by  igniting  the  double  ammonium 
phosphate  precipitated  from  beryllium  solutions  by  microcosmic 
salt  or  ammonium  phosphate  in  presence  of  ammonium  chloride. 
In  presence  of  the  proper  amount  of  ammonium  chloride  (10 
grm.  to  20  grm.  in  100  cm3-200  cm3  of  liquid)  zinc  ammonium 


ZINC,  AND  CADMIUM  IN  ANALYSIS.  265 

phosphate  can  be  obtained  in  the  ideal  condition,  which  on 
ignition  yields  the  pyrophosphate.  This  method  may  serve, 
therefore,  for  the  accurate  estimation  of  zinc. 

Cadmium  may  be  estimated  with  accuracy  as  the  pyrophos- 
phate if  the  precipitate  by  microcosmic  salt  in  the  nearly 
neutral  solution  containing  ammonium  chloride  in  the 
proportion  of  ten  grams  to  one  hundred  cubic  centimeters  is 
allowed  to  stand  several  hours  before  filtering.  In  this  way 
all  cadmium  separates  out  from  the  solution  as  a  beautiful 
crystalline  mass  of  cadmium  ammonium  phosphate  of  ideal 
constitution.  The  conditions,  must,  however,  be  preserved 
with  care ;  there  must  be  no  excess  of  ammonia,  no  free  acid, 
and  no  excess  of  ammonium  salt  beyond  the  quantity  indicated, 
while  that  amount  is  necessary. 


XXXII 

SEPARATION  OF  IRON  FROM  CHROMIUM,  ZIRCONIUM, 
AND  BERYLLIUM,  BY  THE  ACTION  OF  GASEOUS 
HYDROCHLORIC  ACID  ON  THE  OXIDES. 

BY  FRANKE  STUART  HAVENS  AND  ARTHUR  FITCH  WAY.* 

IT  has  been  shown  in  a  former  paper  from  this  laboratory  f 
that  iron  oxide  may  be  completely  volatilized  as  chloride 
by  a  strong  current  of  hydrochloric  acid  gas  acting  at  a 
temperature  of  450  -500°,  and  also  that  the  addition  of  a  little 
free  chlorine  to  the  gaseous  hydrochloric  acid  renders  this 
action  complete  at  lower  temperatures,  180°-200°,  without  the 
danger  of  error  arising  from  the  liability  of  ferric  chloride  to 
dissociation,  or  from  deficiency  of  oxidation  in  the  oxide 
treated,  or  mechanical  loss  due  to  too  rapid  volatilization. 
It  has  also  been  shown  that  this  reaction  can  be  employed  for 
the  separation  of  iron  and  aluminum,  taken  as  the  oxides,  and 
its  application  to  the  separation  of  iron  from  other  metallic 
oxides  has  been  suggested. 

The  oxides  of  chromium,  zirconium,  and  beryllium,  like 
aluminum  oxide,  are  not  acted  upon  by  a  current  of  dry 
hydrochloric  acid  gas  at  the  temperatures  before  mentioned, 
and  these  oxides  also  can  be  entirely  freed  from  iron  by  this 
reaction,  as  the  experiments  to  be  described  will  show.  The 
procedure  was  the  same  in  each  case  and  analogous  to 
that  employed  for  the  separation  of  iron  from  aluminum.  A 
mixture  of  a  weighed  portion  of  one  of  these  oxides  with  a 
weighed  portion  of  ferric  oxide,  contained  in  a  porcelain  boat 
and  placed  within  a  roomy  glass  tube  supported  in  a  small 

*  From  Am.  Jour.  Sci.,  viii,  217. 

t  Gooch  and  Havens,  Am.  Jour.  Sci.,  vii,  370.    This  volume,  p.  215. 


SEPARATION  OF  IRON  FROM  CHROMIUM,  ETC.       267 


combustion  furnace,  was  submitted  to  the  action  of  a  dry 
current  of  hydrochloric  acid  gas  and  chlorine  generated  by 
dropping  sulphuric  acid  upon  a  mixture  of  strong  hydrochloric 
acid,  common  salt,  and  a  small  amount  of  manganese  dioxide. 
The  gas  was  admitted  at  one  end  of  the  combustion  tube  and 
passed  out  at  the  other  through  a  water  trap,  while  the  required 
temperature,  from  200° -300°,  was  maintained  by  regulating 
the  various  burners  of  the  furnace.  The  time  of  action  varies 
somewhat  with  the  condition  of  the  oxide  to  be  volatilized,  and 
the  temperature;  generally  an  hour's  heating  at  200°,  proves 
sufficient  for  the  complete  removal  of  0.1  grm.  of  iron.  At 
higher  temperatures  the  action  is  more  rapid ;  but  the  lighter 
oxide,  the  beryllium  especially,  is  liable  to  mechanical  loss 
through  the  too  rapid  volatilization  of  the  iron,  as  experiment 
(17),  where  a  temperature  of  500°  was  used,  will  show.  It  is 
better,  therefore,  to  use  lower  temperatures,  raising  the  heat 
for  a  few  minutes  when  the  action  is  apparently  complete  to 
ensure  the  removal  of  the  last  traces  of  iron.  Tests  showed 


Exp. 

Fe203  taken. 

Cr,O3  taken. 

Cr2O8  found. 

Error. 

(1) 

(2) 
(8) 
(4) 
(5) 
(6) 

grm. 

0.1007 
0.1007 
0.1010 
0.1019 
0.2007 

grm. 
0.1008 
0.1006 
0.1000 
0.1005 
0.1006 
0.1003 

grm. 
0.1008 
0.1006 
0.1002 
0.1003 
0.1005 
0.0999 

grm. 

0.0000 
0.0000 
0.0002+ 
0.0002- 
0.0001- 
0.0004- 

ZrO,  taken. 

ZrO,  found. 

(7) 
(8) 
(9) 
(10) 

(11) 

0.1053 
0.1204 
0.1236 
0.2150 

0.1516 
0.1010 
0.1519 
0.1516 
0.1517 

0.1516 
0.1010 
0.1523 
0.1517 
0.1519 

0.0000 
0.0000 
0.0004+ 
0.0001+ 
0.0002+ 

BeO  taken. 

BeO  found. 

(12) 
(13) 
(14) 
(16) 
(16) 
(17) 
18) 

0.0997 
0.1045 
0.1215 
0.1510 
0.0230 

0.1309 
0.1285 
0.0456 
0.1099 
0.1080 
0.1305 
0.1081 

0.1311 
0.1285 
0.0457 
0.1099 
0.1081 
0.1290 
0.1083 

0.0002+ 
0.0000 
0.0001+ 
0.0000 
0.0001+ 
0.0015- 
0.0002+ 

268       SEPARATION  OF  IRON  FROM  CHROMIUM,  ETC. 

the  residual   oxides  from  which  the  ferric  oxide  had  been 
removed  in  this  manner  to  be  entirely  free  from  iron. 

The  separation  of  iron  from  chromium,  zirconium,  and 
beryllium  by  this  method  is  obviously  complete  within  very 
satisfactory  limits  of  error. 


XXXIII 

THE   IODOMETRIC    DETERMINATION    OF   GOLD. 

BY  F.  A  GOOCH  AND  FREDERICK  H.  MORLEY.* 

IN  a  recent  attempt  to  measure  small  amounts  of  gold  in 
solution  by  titrating  with  sodium  thiosulphate  the  iodine  set 
free  in  the  action  of  an  excess  of  potassium  iodide  upon  auric 
chloride,  Petersonf  has  been  led  to  conclude  that,  on  the  aver- 
age, one-half  more  thiosulphate  is  used  up  in  changing  the 
characteristic  starch  iodide  blue  to  the  faint  rose  color  which 
precedes  entire  bleaching  than  is  called  for  upon  the  theory 
that  the  thiosulphate  is  simply  converted  to  the  tetrathionate 
in  the  usual  manner.  Peterson  explains  the  anomaly  upon  the 
hypothesis  that,  besides  acting  upon  the  free  iodine,  the  thio- 
sulphate is  used  up  coincidently  by  interaction  with  the  aurous 
salt,  formed  in  the  reduction,  with  formation  of  a  gold  sodium 
thiosulphate  on  the  type  of  the  well-known  silver  sodium  thio- 
sulphate. The  reaction  of  this  hypothesis  is  in  the  nature  of 
things  most  improbable,  since  there  is  no  reason  to  suppose 
that  the  soluble  double  thiosulphate  could  resist  the  action  of 
the  free  iodine  which  is  present  to  the  end  —  the  appearance 
of  the  rose  color,  and  our  study  of  the  reaction  of  sodium  thio- 
sulphate upon  the  mixture  of  gold  chloride  and  potassium 
iodide,  the  account  of  which  follows,  discloses  no  evidence  of 
the  consumption  of  more  thiosulphate  than  is  demanded  by 
the  usual  theory,  which  postulates  the  simple  formation  of  the 
tetrathionate  by  the  interaction  of  the  thiosulphate  and  free 
iodine. 

It  appeared  in  the  course  of  our  preliminary  experimenta- 

*  From  Am.  Jour.  Sci.,  viii,  261. 
t  Zeitschr.  anorg.  Chem.,  xix,  63. 


270 


IODOMETRIC  DETERMINATION  OF  GOLD. 


tion  that,  while  practically  similar  results  were  obtained  by 
adding  the  thiosulphate  until  the  blue  of  the  starch  iodide  had 
changed  to  rose,  the  indications  were  somewhat  more  con- 
cordant when  the  final  rose  color  was  developed  by  adding 
iodine  to  the  solution  from  which  the  blue  had  been  bleached 
to  colorlessness  by  a  slight  excess  of  the  thiosulphate. 

It  appeared,  also,  that  the  reduction  of  the  auric  salt,  with 
the  consequent  liberation  of  iodine,  is  conditioned  by  the  vol- 
ume of  the  solution,  the  mass  of  the  iodine  present,  and  the 
time  of  action. 

The  following  statement,  in  which  each  result  is  the  average 
of  several  titrations  in  close  agreement,  shows  the  effect  upon 
the  immediate  evolution  of  iodine  brought  about  by  adding 
varying  amounts  of  water  to  the  gold  solution  before  introduc- 
ing the  iodide,  and  the  effect  of  different  amounts  of  iodide 
at  different  dilutions. 


Potassium  iodide. 

Gold 
chloride. 

Volume 
before  the 

addition  of 

the  thio- 

0.01 grm. 

0.02  grm. 

0.05  grm. 

0.1  grm. 

0.2  grm. 

0.00087  grm. 

sulphate. 

.sF  . 

fO.81 

0.81 

0.81 

0.82 

0.84 

0.00087 

cm3 
15 

If1?"!. 

0.77 
0.74 

0.78 
0.72 

0.80 
0.78 

0.81 
0.79 

0.81 

0.80 

0.00087 
0.00087 

25 
60 

11  1  S 

0.61 

0.61 

0.68 

0.76 

0.79 

0.00087 

100 

2  fl 

[0.45 

0.49 

0.60 

0.72 

0.75 

0.00087 

200 

It  is  evident  that  for  the  smaller  amounts  of  iodide  the  lib- 
eration of  iodine  decreases  rapidly  with  the  dilution.  The 
larger  amounts  at  the  highest  concentration  show  readings 
a  trifle  above  the  normal — perhaps  because  the  well-known 
effect  of  concentrated  solutions  of  a  soluble  iodide  upon  the 
delicacy  of  the  starch  end-color  begins  to  appear.  At  vol- 
umes lying  between  the  limit  of  25  cm3  and  50  cm3  0.1 
grm.  of  potassium  iodide  is  an  appropriate  amount  to  use; 
at  a  volume  of  15  cm3,  0.01  grm.  to  0.05  grm.  of  the  iodide 
will  do  the  work;  and  at  lower  dilutions,  as  will  appear  in 


IODOMETRIC  DETERMINATION  OF  GOLD.         271 

the  tabular  statements  to  follow,  even  less  of  the  iodide  is 
effective. 

In  the  series  of  experiments  of  which  the  details  are  given 
in  Table  I,  use  was  made  of  a  solution  of  pure  gold  chloride 
containing  0.8710  grm.  to  the  liter  —  as  determined  by  careful 
precipitation  in  the  usual  manner  by  ferrous  sulphate,  and  by 
an  alkaline  solution  of  formaldehyde  according  to  the  method 
of  Vanino.*  A  nearly  centinormal  solution  of  iodine  was  pre- 
pared by  diluting  to  a  liter  100  cm3  of  nearly  decinormal  iodine 
in  potassium  iodide  carefully  standardized  against  exactly 
decinormal  arsenious  acid.  A  nearly  centinormal  solution  of 
sodium  thiosulphate  (containing  1.7012  grm.  of  Na2S2O8  to 
the  liter)  was  made  by  diluting  to  a  liter  100  cm3  of  a  nearly 
decinormal  solution  of  that  reagent  which  had  been  standard- 
ized carefully  against  the  standard  iodine  prepared  as  described. 
The  solution  of  potassium  iodide  employed  contained  10  grm. 
of  that  salt  in  the  liter. 

In  conducting  the  experiments,  a  convenient  amount  of  the 
solution  of  gold  chloride  was  drawn  from  a  burette,  potassium 
iodide  was  introduced  in  the  amounts  indicated  (always  several 
times  the  theoretical  equivalent  of  the  gold,  and  more  than 
enough  to  dissolve  the  aurous  iodide  precipitated  at  first),  a 
sufficiency  of  clear  starch  indicator  was  added,  the  starch  blue 
was  bleached  by  the  thiosulphate,  and  the  iodine  was  added 
until  the  liquid  assumed  a  faint  rose  color.  Upon  the  theory 
that  potassium  iodide  sets  free  two  atoms  of  iodine  for  every 
molecule  of  auric  chloride  (or  every  atom  of  gold)  present, 
and  that  the  thiosulphate  acts  only  upon  the  free  iodine  to 
form  the  tetrathionate  in  the  usual  manner,  every  cubic  centi- 
meter of  the  thiosulphate  solution  used  in  the  reaction  after 
deducting  the  amount  equivalent  to  the  iodine  introduced  to 
get  the  end-color,  should  represent 

1Q7  3 

X  0.0017012  =  0.001061  grm.  of  gold. 


2(158.22) 

»  Ber.  Dtsch.  chem.  Ges.,  xxi,  1763. 


272         IODOMETRIC  DETERMINATION  OF  GOLD. 

TABLE  I. 


Gold  chloride                              =  0.8710  to  1  liter. 

Sodium  thiosulphate,  nearly  ~  =  1.7012    "      " 

Iodine,  nearly  ^                        =1.3697    "      " 

Volume  at  beginning  of  titration,  approximately  50  cm8. 

AuCl3 
taken. 

KI 

taken. 

Na2S208 
used. 

Gold  found. 

Theory 
for  gold. 

Error. 

Per  cent. 

cm3 

grm. 

cm3 

grm. 

grm. 

grin. 

(1) 

5 

0.05 

4.02 

0.00426 

0.00435 

0.00009- 

2.1 

(2) 

6 

0.05 

4.01 

0.00425 

0.00435 

0.00010- 

2.3 

3) 

5 

0.05 

4.06 

0.00431 

0.00435 

0.00004- 

0.9 

4) 

5 

0.05 

4.07 

0.00432 

0.00435 

0.00003- 

0.7 

5) 

5 

0.05 

4.04 

0.00428 

0.00435 

0.00007- 

1.6 

6) 

10 

0.08 

8.17 

0.00867 

0.00871 

0.00004- 

0.5 

7) 

10 

0.08 

8.15 

0.00864 

0.00871 

0.00007- 

0.8 

8) 

10 

0.08 

8.16 

0.00865 

0.00871 

0.00006- 

0.7 

9 

10 

0.08 

8.15 

0.00864 

0.00871 

0.00007- 

0.8 

(10) 

10 

0.08 

8.19 

0.00869 

0.00871 

0.00002- 

0.2 

(11) 

10 

0.08 

8.46 

0.00897 

0.00871 

0.00026+ 

3.0 

(12) 

10 

0.08 

8.24 

0.00874 

0.00871 

0.00003+ 

0.3 

Plainly,  these  results  accord  reasonably  with  the  theory  that 
two  molecules  of  the  thiosulphate  are  the  equivalent  in  this 
reaction  of  two  atoms  of  iodine  and  one  atom  of  gold.  There 
is  no  evidence  whatever  of  the  excessive  action  affirmed  by 
Peterson. 

The  strength  of  the  standard  solutions  used  in  the  experi- 
ments described  was  such  that  an  error  of  0.01  cm3  in  reading 
the  volumes  used  would  correspond  to  an  error  of  0.00001 
grm.  of  gold.  It  is  not  to  be  expected  that  such  readings  can 
be  trusted  ordinarily  to  a  higher  degree  of  accuracy  than  0.02 
cm3.  In  case  all  three  solutions  should  be  read  to  this  limit 
of  accuracy  with  the  errors  of  all  lying  in  the  same  direction, 
the  summation  of  error  would  correspond  to  0.00006  grm.  of 
gold. 

In  the  following  experiments,  therefore,  solutions  obtained 
by  properly  diluting  those  of  the  previous  series  were  em- 
ployed. The  use  of  a  more  dilute  solution  of  gold  obviated 
the  necessity  for  diluting  the  mixture  of  gold  chloride  and  the 
iodide  before  titrating  with  the  thiosulphate.  It  was  found, 


IODOMETRIC  DETERMINATION  OF  GOLD.         273 
TABLE  II. 


A. 

Gold  chloride                              =  0.0871  to  1  liter. 

Sodium  thiosulphate,  nearly  ^  =  1.7012    "     " 

Iodine,  nearly  j£                        =1.3697    "     « 

Solution  of  gold  chloride  not  diluted  before  mixing  with  potassium  iodide. 

Exp. 

AuCls 
taken. 

Kl 

taken. 

JSra,BJO, 
used. 

Gold 
taken. 

Gold 
found. 

Error. 

cm8 

grin. 

cm» 

grm. 

grm. 

grm. 

(1) 

10 

0.01 

0.83 

0.00087 

0.00088 

0.00001+ 

2) 

10 

0.01 

0.83 

0.00087 

0.00088 

0.00001+ 

3) 

10 

0.01 

0.80 

0.00087 

0.00085 

0.00002- 

4) 

10 

0.02 

0.84 

0.00087 

0.00089 

0.00002+ 

5) 

10 

0.02 

0.88 

0.00087 

0.00093 

0.00006+ 

6) 

10 

0.02 

0.82 

0.00087 

0.00087 

0.00000 

(7) 

10 

0.02 

0.88 

0.00087 

0.00093 

0.00006+ 

(8) 

10 

0.02 

0.83 

0.00087 

0.00088 

0.00001+ 

(9) 

10 

0.10 

0.80 

0.00087 

0.00085 

0.00002- 

(10) 

10 

0.10 

0.82 

0.00087 

0.00087 

0.00000 

(11 

10 

0.01 

0.83 

0.00087 

0.00088 

0.00001+ 

(12) 

9 

0.01 

0.73 

0.00078 

0.00077 

0.00001- 

(13) 

8 

0.01 

0.65 

0.00070 

0.00069 

0.00001- 

(14) 

7 

0.01 

0.58 

0.00061 

0.00061 

0.00000 

(15) 

6 

0.008 

0.51 

0.00052 

0.00054 

0.00002+ 

(16) 

6 

0.008 

0.41 

0.00043 

0.00044 

0.00001+ 

(17) 
(18) 

4 
3 

0.005 
0.005 

0.35 
0/24 

0.00035 
0.00026 

0.00037 
0.00026 

0.00002+ 
0.00000 

(19) 

2 

0.003 

0.21 

0.00017 

0.00022 

0.00005+ 

(20) 

1 

0.003 

0.10 

0.00009 

0.00011 

0.00002+ 

B. 

Gold  chloride                              =0.0871    to  1  liter. 

Sodium  thiosulphate,  nearly  ~  =  0.17012    "     " 

Iodine,  nearly  ^                         =0.13697    "     " 

(21) 

10 

0.01 

8.39 

0.000871 

0.000890 

0.000019+ 

(22) 

9 

0.01 

7.45 

0.000784 

0.000790 

0.000006+ 

(23) 

8 

0.01 

6.30 

0.000697 

0.000668 

0.000029- 

(24) 

7 

0.008 

5.50 

0.000610 

0.000583 

0.000027- 

(25) 

6 

0.008 

5.12 

0.000523 

0.000543 

0.000020+ 

(26) 

5 

0.005 

4.23 

0.000435 

0.000449 

0.000014+ 

(27) 

4 

0.005 

3.38 

0.000348 

0.000358 

0.000010+ 

(28) 

3 

0.003 

2.55 

0.000261 

0.000270 

0.000009+ 

(29) 

2 

0.003 

1.71 

0.000174 

0.000181 

0.000007+ 

(30) 

1 

0.003 

0.90 

0.000087 

0.000095 

0.000008+ 

however,  that  when  the  T^-  solution  of  iodine  is  employed  a 
correction  of  0.1  cm3  for  volumes  not  exceeding  30  cm3  be- 


VOL.    II. 


18 


274         IODOMETRIC  DETERMINATION  OF  GOLD. 

comes  necessary  —  the  amount  required  to  bring  out  the  rose 
color  in  blank  tests  containing  no  gold.  After  the  introduc- 
tion of  0.1  cm3  of  1-jW5  iodine  into  a  mixture  of  potassium 
iodide  and  starch  indicator  of  volume  not  exceeding  30  cm3,  a 
single  drop  of  the  gold  solution  —  equivalent  to  0.000002  grm. 
of  gold  —  gave  a  distinct  rose  color :  before  such  adjustment 
of  the  solution  five  drops  —  equivalent  to  0.000010  of  gold  — 
were  needed  to  develop  the  same  color. 

These  results  run  on  the  whole  as  regularly  as  could  be 
expected,  and  the  use  of  the  dilute  standard  solutions  is 
obviously  of  advantage. 

In  the  practical  application  of  any  such  process  for  the 
determination  of  gold,  the  elementary  form  of  that  metal  is 
the  natural  starting-point.  To  get  the  metal  into  solution 
with  chlorine  water  or  mixed  hydrochloric  and  nitric  acids  is 
an  easy  matter,  but  the  removal  of  the  excess  of  the  oxidizer 
by  evaporation  without  reducing  some  auric  chloride  to  the 
aurous  form  is  difficult.  We  have  found,  however,  that  the 
free  chlorine  may  be  removed  from  a  solution  of  auric 
chloride,  without  reducing  the  auric  salt,  by  treatment  of  the 
solution  with  ammonia  in  excess,  boiling  gently,  acidifying 
with  hydrochloric  acid  and  heating  if  necessary  to  redissolve 
the  precipitate  by  ammonia,  again  treating  with  ammonia  and 
heating,  and  once  more  acidifying.  On  the  second  addition 
of  ammonia  no  precipitation  usually  takes  place  with  the 
amounts  of  gold  which  we  have  thus  handled,  perhaps  because 
enough  ammonium  chloride  has  been  found  to  hold  it  up. 

The  following  table  contains  determinations  made  with 
such  a  solution  of  pure  gold  leaf  —  tested  gravimetrically  as 
to  purity. 

Obviously,  this  method,  which  rests  upon  the  hypothesis 
that  sodium  thiosulphate  acts  in  the  normal  manner  only 
upon  the  iodine  set  free  by  the  interaction  of  gold  chloride 
and  potassium  iodide,  offers  trustworthy  means  for  the 
determination  of  small  amounts  of  gold. 


IODOMETRIC  DETERMINATION  OF  GOLD.         275 
TABLE  III. 


Gold  chloride  made  by  dissolving  0.0104  grm.  of  pure  gold  in  the  manner 

described  and  diluting  to  200  cm8. 

Sodium  thiosulphate,  nearly  j^  =  0.17012  to  1  liter. 

Iodine,  nearly  ~                         =0.13697    "      " 

Potassium  iodide                          =  10  grm.    "      " 

Portions  were  treated  with  the  potassium  iodide  without  previous  dilution. 

Eip. 

AuCl. 
taken. 

KI 

taken. 

«&°' 

Gold 
taken. 

Gold 
found. 

Error. 

cm3 

grm. 

cm** 

grm. 

grm. 

gnn. 

(1) 

1 

0.005 

0.55 

0.000052 

0.000058 

0.000006+ 

(2) 

1 

0.005 

0.55 

0.000052 

0.000058 

0.000006+ 

(3) 

2 

0.005 

1.06 

0.000104 

0.000112 

0.000008+ 

(4) 

2 

0.005 

1.08 

0.000104 

0.000114 

0.000010+ 

(5) 

5 

0.01 

2.45 

0.000260 

0.000260 

0.000000 

(6 

5 

0.01 

2.50 

0.000260 

0.000265 

0.000005+ 

7) 

5 

0.01 

2.45 

0.000260 

0.000260 

0.000000 

(8) 

5 

0.01 

2.50 

0.000260 

0.000265 

0.000005+ 

(9) 

5 

0.01 

2.50 

0.000260 

0.000265 

0.000005+ 

(10) 

10 

0.02 

4.86 

0.000520 

0.000515 

0.000005- 

(11) 

10 

0.02 

4.85 

0.000520 

0.000517 

0.000003- 

(12) 

10 

0.02 

4.90 

0.000520 

0.000520 

0.000000 

(13) 

10 

0.02 

4.80 

0.000520 

0.000512 

0.000008- 

(14) 

10 

0.02 

4.84 

0.000520 

0.000516 

0.000004- 

XXXIV 

THE  ACTION  OF  ACETYLENE  ON  THE  OXIDES 
OF  COPPER. 

BY  F.  A.  GOOCH  AND  DEFOREST  BALDWIN.* 

IN  a  recent  paper  by  Erdmann  and  Kothnerf  an  account  is 
given  of  the  formation  of  a  peculiar,  light-brown,  highly 
voluminous  substance  by  the  action  of  acetylene  below  250° 
C.  upon  cuprous  oxide,  or  even  (though  more  slowly)  upon 
copper.  The  product  obtained  by  passing  acetylene  during 
eighteen  hours  over  1  grm.  of  cuprous  oxide  (prepared 
from  copper  sulphate,  grape  sugar,  and  sodium  hydroxide) 
amounted  to  7  grm.  and  filled  a  space  of  nearly  300  cm3. 
At  higher  temperatures  a  black  carbonaceous  mass  is  the 
result,  and  at  red  heat  (400°-500°  C.)  carbon  is  deposited  in 
graphitic  condition.  The  light-brown  fluffy  material  yielded 
cuprous  chloride  to  hydrochloric  acid,  a  distillate  from  its 
mixture  with  zinc  dust  possessing  the  characteristics  of 
naphthene  or,  at  higher  temperature  and  under  rapid  heating, 
aromatic  compounds  among  which  naphthalene  and  a  kresol 
were  indicated.  Erdmann  and  Kothner  classify  this  body  as 
a  very  complex  but  non-explosive  copper  acetylene  (acetylen- 
kupfer,),  and  from  their  analyses  deduce  the  formula  C44H64 
Cus.  Apart  from  the  unusual  constitution  of  this  symbol, 
its  most  striking  peculiarity  is  that  it  implies  a  loss  of  carbon, 
rather  than  hydrogen,  from  the  acetylene  in  the  reaction  with 
cuprous  oxide — a  condition  of  affairs  which  would  be  most 
remarkable  in  the  light  of  Campbell's  experience,^  according 
to  which  acetylene  passed  over  palladinized  copper  oxide 

*  From  Am.  Jour.  Sci.,  viii,  354.         t  Zeitschr.  anorg.  Chem.,  xviii,  49. 
t  Amer.  Chem.  Jour.,  xvii,  690. 


ACTION  OF  ACETYLENE  ON  OXIDES  OF  COPPER.    277 

yielded  water  at  225°-230°  and  carbon  dioxide  only  when 
the  temperature  rose  to  315°-320°  with  the  formation  of  a 
black  deposit.  Upon  scrutinizing  the  figures  of  Erdmann  and 
Kbthner  with  care,  however,  it  appears  that  the  formula 
given  by  these  investigators  rests  upon  some  oversight  in 
calculation:  the  ratio  of  carbon  atoms  to  hydrogen  atoms 
proves  to  be  actually,  according  to  the  data  given,  6.45  :  5.70 ; 
which  means,  of  course,  that  the  new  product  is  deficient,  as 
would  be  expected,  in  hydrogen  (not  in  carbon)  as  compared 
with  acetylene. 

As  to  the  content  of  the  new  substance  in  copper,  the 
analytical  data  are  unfortunately  ambiguous;  for  we  note 
the  weights  found  of  copper  oxide  converted  into  percentages 
of  copper  without  preliminary  reduction.  If  the  fault  is 
typographical  and  in  the  analytical  data,  the  calculated 
percentages  of  copper  being  correct,  the  average  percentage 
of  copper  amounts  to  15.43 :  if,  on  the  other  hand,  the  ana- 
lytical data  are  right,  the  error  being  in  their  reduction,  the 
percentage  of  copper  amounts  to  12.92.  In  the  one  case  the 
summation  of  the  analysis  leaves  a  deficiency  of  about  1.5 
per  cent,  and  in  the  other  of  about  4  per  cent,  which  hi  either 
case  may  really  represent  oxygen  hi  the  substance.  This 
condition  of  matters  leaves  the  "  acetylen-kupfer "  of  Erd- 
mann and  Kbthner  in  uncertain  standing. 

More  than  thirty  years  ago  it  was  noticed  by  Berthelot  * 
that  acetylene  is  polymerized  by  heat  or  decomposed  partially 
into  carbon  and  hydrogen,  and  that  such  action  takes  place 
more  readily  and  at  lower  temperatures  in  presence  of  metallic 
iron  with  production  of  carbon,  hydrogen  and  compounds 
different  from  those  formed  by  heat  alone. 

Moissan  and  Moureu  f  have  observed  the  incandescence  of 
acetylene  passed  over  finely  divided  iron,  cobalt,  nickel,  or 
platinum  at  the  ordinary  temperature,  with  production  of 
carbon,  hydrogen,  and  pyrogenic  compounds,  and  have  found 
the  occasion  of  such  behavior  in  the  porosity  of  the  metals 
employed. 

*  Ann.  Chim.  [4],  ix,  448.  t  Compt.  rend.,  cxxii,  1240. 


278  THE  ACTION  OF  ACETYLENE 

It  would  seem  natural,  however,  that  the  presence  of  oxygen, 
free  or  combined,  may  also  play  a  considerable  part  in  such 
phenomena,  just  as  appears  to  be  the  case  in  the  peculiar  action 
recorded  by  Gruner  J  of  carbon  monoxide  upon  iron  reduced 
by  hydrogen,  which,  as  Moissan  has  shown,§  is  produced  pure 
only  with  the  greatest  precaution  and  generally  carries  a 
large  proportion  of  ferrous  oxide.  The  fact  that  the  "  acety- 
len-kupfer  "  of  Erdmann  and  Kothner  is  produced  more  easily 
by  the  action  of  cuprous  oxide  upon  acetylene  than  by  the 
action  of  metallic  copper  upon  acetylene,  suggests  that  it  may 
be  the  oxidizing  power  of  the  cuprous  oxide  which  gives 
to  this  reagent  its  peculiar  activity.  The  question  arises, 
therefore,  as  to  whether  the  copper  is  in  reality  an  essential 
constituent  of  the  compound  of  Erdmann  and  Kothner. 

In  our  experiments  upon  the  action  of  acetylene  upon  the 
oxides  of  copper  (and  other  elements)  we  have  conducted  the 
gas  (made  in  the  ordinary  way  by  the  action  of  water  on 
calcium  carbide,  and  kept  over  water)  over  the  oxide  contained 
in  a  porcelain  boat  placed  within  a  glass  tube,  2  cm.  in  diameter 
and  50  cm.  long,  which  was  heated  over  a  small  combustion 
furnace.  The  glass  tube  was  fitted  at  each  end  with  a  rubber 
stopper,  one  carrying  a  smaller  tube  for  the  introduction  of  the 
acetylene  and  a  high-temperature  thermometer  so  held  that  its 
bulb  rested  horizontally  immediately  over  the  boat  containing 
the  oxide,  while  the  other  was  fitted  with  a  water-trap.  In  the 
preh'minary  experiments  no  attempt  was  made  to  purify  the 
acetylene  employed  other  than  to  keep  it  over  water,  or,  since 
water  is  a  product  of  its  action  upon  oxides,  to  dry  it :  in  later 
experiments  to  secure  products  for  careful  analysis  it  was  dried 
and  purified  with  care. 

We  found  that  225°  C.  is  the  temperature  most  favorable 
for  the  formation  of  the  voluminous  product  obtained  by  acting 
with  acetylene  upon  cuprous  oxide  as  described  by  Erdmann 
and  Kothner.  At  this  temperature  the  tube  is  choked  rapidly 
with  the  fluffy  product  and  water  forms,  but,  as  Campbell 
found  in  his  experiments  upon  palladinized  copper  oxide,  no 

t  Ann.  Chim.  [4],  xxvi,  6.  §  Ann.  Chim.  [5],  xxi,  199. 


ON  THE  OXIDES  OF  COPPER.  279 

appreciable  amount  of  carbon  dioxide  is  produced.  The 
content  of  the  product  in  copper  varies  in  the  sample  and  in 
different  experiments,  our  results  lying  between  1.54  per  cent 
and  24.21  per  cent  of  the  substance  taken  for  ignition. 

It  appeared,  also,  that  the  action  of  acetylene  upon  cupric 
oxide  is  precisely  similar  to  that  upon  cuprous  oxide  excepting 
the  evident  reduction  of  the  former  oxide  early  in  the  action. 
The  amount  of  copper  in  the  product  of  such  action  varied  in 
our  experiments  from  6.53  per  cent  to  21.30  per  cent.  In  one 
case  the  experiment  of  re-submitting  to  the  action  of  acetylene 
a  product  containing  9.34  per  cent  of  copper  was  made  with 
the  result  that  a  new  growth  of  the  substance  formed  which 
on  analysis  yielded  3.87  per  cent  of  copper. 

A  roll  of  copper  gauze  carefully  reduced  in  hydrogen  and 
then  oxidized  at  one  end  in  the  outer  flame  of  a  Bunsen 
burner  gave,  when  acted  upon  by  acetylene  at  225°-250°  C., 
the  characteristic  deposit  upon  the  oxidized  end  only,  the 
unoxidized  end  being  merely  discolored. 

These  results  go  to  show  that,  while  metallic  copper  may  at 
comparatively  high  temperatures  induce  the  polymerization  of 
acetylene,  it  is  an  oxidiiing  action  which  starts  at  moderately 
low  temperatures  the  formation  of  the  peculiar  derivatives 
under  consideration.  Thus  we  find  that  ferric  oxide  heated  in 
acetylene  at  temperatures  varying  from  150°  to  360°,  accord- 
ing to  circumstances,  darkens,  glows,  and  gathers  with  evolution 
of  heat  a  dark  carbonaceous  deposit.  In  the  products  of  such 
action  we  have  found  the  content  of  iron  varying  from  2.80 
per  cent  to  5.86  per  cent. 

Silver  oxide,  too,  acts  upon  acetylene:  thus,  in  one 
experiment,  action  was  evident  at  the  ordinary  atmospheric 
temperature,  and  a  violent  explosion,  which  completely  shat- 
tered the  boat  and  scattered  metallic  silver  upon  the  sides  of 
the  glass  tube,  followed  before  the  temperature  reached  100°. 

In  the  locally  violent  explosion  of  the  last  experiment 
we  have  evidence  of  the  formation  in  the  early  stage  of  an 
acetylide  which  is  decomposed  later  when  the  temperature  of 
dissociation  is  reached.  In  the  experiments  with  the  oxides  of 


280 


THE  ACTION  OF  ACETYLENE 


copper  and  iron  the  temperature  at  which  the  acetylene  begins 
to  act  is  evidently  above  the  point  at  which  sensitive  acetylides 
would  naturally  dissociate,  and  we  have  in  the  observed 
phenomena  no  evidence  of  the  formation  of  such  compounds  of 
copper  and  iron  under  the  conditions  of  experimentation. 

In  experiments  (1)  to  (3)  of  the  following  table  are  given 
the  results  of  the  analysis  of  several  products  obtained  by 
conducting  acetylene  (purified  by  passing  through  a  solution 
of  mercuric  chloride  in  hydrochloric  acid  and  dried  over  caustic 
potash)  over  pure  cuprous  oxide.  The  temperature  was  kept 
in  these  experiments  at  225°,  and  in  the  course  of  a  half-hour 
the  tube  was  choked  completely  by  material  compacted  by 
the  pressure  to  (1)  a  spongy  mass  of  light-brown  color  on  the 
exterior  next  the  walls  of  the  tube,  (2)  darker  within  and  (3) 
nearly  black  in  the  bottom  of  the  boat,  where  the  cuprous 
oxide  lay  originally. 


Ezp. 

Weight 
of  sub- 
stance 
taken. 

Found. 

Calculated. 

CO, 

H2O 

CuO 

C 

H 

Cu 

Oby 
differ- 
ence. 

(1) 
(2) 
(3) 

grin. 

0.1170 
0.2247 
0.1096 

grnit 

0.3978 
0.7489 
0.3678 

gnu. 
0.0673 
0.0979 
0.0488 

gnu. 
0.0022 

o!o045 

gnu. 
0.1085 
0.2042 
0.1003 

grm. 
0.0075 
0.0109 
0.0054 

grm. 
0.0018 

grm. 

0.0008- 

0.0036 

0.0003 

(4) 
(5) 

0.1360 
0.1188 

0.4116 
0.3098 

0.0579 
0.0461 

0.0182 
0.0317 

0.1123 
0.0845 

0.0064 
0.0051 

0.0146 
0.0253 

0.0027 
0.0039 

Per  cent  of  carbon 
Pet  cent  of  hydrogen 
Per  cent  of  copper 
Per  cent  of  oxygen 


(1) 
92.74 

(2) 
90.88 

(3) 
91.51 

(4) 

82.57 

(6) 
71.13 

6.41 

4.85 

4.93 

4.71 

4.29 

1.54 

3.29 
0.27 

10.74 
1.98 

21.30 
3.20 

100.69 


100.00     100.00     100.00 


In  experiments  (4)  and  (5)  the  substances  analyzed  repre- 
sent the  products  of  the  action  of  acetylene  (not  specially 
purified)  on  cupric  oxide. 

The  oxygen  present  in  these  products  is  obviously  pro- 
portional to  the  amount  of  copper  and  is  never  more  than 


ON  THE  OXIDES  OF  COPPER.         281 

enough  to  be  completely  accounted  for  upon  the  supposition 
that  some  of  the  original  oxide  taken  stills  holds  its  oxygen. 
So  far  as  the  analyses  show,  the  product  of  lightest  color  (1) 
contains  very  little  copper  and  no  oxygen ;  the  darkest  prod- 
uct (3)  obtained  from  the  cuprous  oxide  contains  oxygen 
corresponding  to  a  mixture  of  two  parts  of  copper  with  three 
parts  of  cuprous  oxide;  the  oxygen  in  the  products  of  (4) 
and  (5)  obtained  by  acting  upon  cupric  oxide  is  approximately 
enough  to  correspond  to  a  mixture  of  cuprous  and  cupric 
oxides  in  equal  proportions.  This  fact,  taken  in  connection 
with  the  great  range  of  variation  in  proportion  and  the 
minimum  to  which  the  copper  falls  in  the  product,  which 
would  be  least  likely  to  include  contaminating  metal  or 
oxide,  suggests  very  strongly  the  probability  that  the  oxygen 
present  is  in  union  with  copper  and  that  the  copper  is  held 
mechanically  as  metal  or  oxide  and  is  not  the  essential 
constituent  of  an  organic  compound.  Leaving  out  of  con- 
sideration, therefore,  the  copper  and  copper  oxides,  and 
calculating  the  composition  of  the  products  assumed  to 
consist  essentially  of  carbon  and  hydrogen,  we  derive  the 
following  statement: 

(1)  (2)  (3)  (4)  (5) 

Per  cent  of  carbon             93.54  94.93  94.88  94.60  94.31 

Per  cent  of  hydrogen           6.46  5.07  5.12  5.40  5.69 

100.00  100.00  100.00  100.00  100.00 

These  figures  correspond  to  symbols  varying  from  Ci2H10  to 
nearly  Ci6Hi0,  with  an  average  approximating  CnHi0,  the 
symbol  of  anthracene  or  paranthracene.  The  analytical  data 
of  Erdmann  and  Kothner  point  in  the  average  to  a  product 
corresponding  more  nearly  to  the  first  of  these  symbols  than 
to  either  of  the  others.  The  product  is  doubtless  variable 
with  the  temperature  and  the  activity  of  oxidation.  Thus, 
in  one  experiment  in  which  acetylene  was  passed  over  ferric 
oxide  the  action  began  at  365°  with  incandescence,  as  de- 
scribed by  Moissan  and  Moureu,*  and  the  analysis  of  the 

*  Loc.  cit 


282    ACTION  OF  ACETYLENE  ON  OXIDES  OF  COPPER. 

product  (carbon  —  91.53,  hydrogen  =  1.36,  Fe  =  5.85,  O  = 
1.26)  indicates  a  proportion  of  carbon  to  hydrogen  about 
four  times  as  great  as  that  of  the  average  product  of  action 
at  225°  on  the  oxides  of  copper. 

Finally,  we  find  no  evidence  that  the  product  of  the  action 
of  acetylene  on  the  oxides  of  copper  under  the  conditions  of 
our  experimentation  is  other  than  a  mixture  of  a  hydro- 
carbon or  hydrocarbons  with  metallic  copper  or  an  oxide  of 
copper,  and,  probably,  in  the  darker  preparations,  some  free 
carbon. 


XXXV 

NOTES  ON  THE  SPACE  ISOMERISM  OF  THE 
TOLUQUINONEOXIME  ETHERS. 

BY  WM.  CONGER  MORGAN.* 

IN  an  article  on  the  "Ethers  of  Toluquinoneoxime  and 
their  Bearing  on  the  Space  Isomerism  of  Nitrogen,"!  pub- 
lished from  this  laboratory,  it  was  stated  that  the  methyl, 
acetyl,  and  benzoyl  ethers  of  toluquinonemetaoxime,  whether 
formed  by  the  action  of  hydroxylamine  on  the  quinone  or  by 
nitrous  acid  on  the  corresponding  cresol,  showed  evidence  of 
existing  in  isomeric  forms.  Of  these  bodies  the  benzoyl  ether 
received  the  most  careful  investigation,  and  by  fractional 
crystallization,  from  the  crude  reaction-product,  two  portions 
were  obtained,  one  readily  separating  from  an  alcoholic  solu- 
tion in  the  form  of  yellow  crystals  melting  at  193°,  while 
a  second  body,  melting  approximately  at  144°,  was  never 
obtained  in  a  state  of  purity.  The  fact  that  from  this 
lower-melting  fraction,  on  recrystallization,  a  portion  of  the 
higher-melting  body  was  always  obtained,  suggested  the 
possibility  that  the  solvent  might  have  a  tendency  to  cause  a 
transition  from  one  isomer  to  the  other.  Since,  however, 
after  repeated  crystallizations  from  boiling  alcohol  a  low- 
melting  fraction  was  obtained,  and  therefore  such  rearrange- 
ment was  evidently  incomplete,  the  action  of  alcohol  under 
pressure  on  the  different  fractions  was  investigated.  From  a 
portion  of  the  ether  melting  at  139°,  heated  for  three  hours  in 
a  closed  tube  to  120°,  the  product  melting  at  193°  crystallized 
in  a  characteristic  form  and  no  crystals  were  obtained  melting 

*  From  Am.  Chem.  Jour.,  xxii,  402. 

t  Bridge  and  Morgan,  Am.  Chem.  Jour.,  xx,  761.    This  volume,  p.  145. 


284  NOTES   ON  THE  SPACE  ISOMERISM 

lower  than  180°.  On  treating  the  higher-melting  isomer  in 
the  same  way,  no  action  was  observed  until  the  tempera- 
ture was  raised  above  150°,  when  complete  decomposition 
ensued. 

In  order  to  ascertain,  if  possible,  a  direct  method  of  tran- 
sition from  one  form  into  the  other,  the  action  of  alkalies, 
among  other  reagents,  was  tried.  No  change  other  than 
a  hydrolytic  cleavage  was  observed,  and  this  was  readily 
completed  on  warming.  By  saponifying  a  fraction  melting 
completely  at  142°  and  treating  the  isolated  oxime  in  the 
usual  way  with  benzoyl  chloride,  the  benzoyl  ether  was  again 
obtained,  liquefying  at  193°,  and,  although  there  was  a 
trifling  irregularity  in  the  melting-point  of  some  of  the  crys- 
tals, no  low-melting  fraction  was  isolated.  On  similar  treat- 
ment of  a  portion  melting  at  193°  the  original  ether  was 
obtained,  but  none  of  the  low-melting  isomer. 

From  these  facts  it  is  obvious  that  the  isomer  existing 
in  a  much  smaller  proportion  in  the  crude  reaction-products, 
present  as  the  principal  constituent  in  the  low-melting  frac- 
tions, must  be  regarded  as  the  labile  form,  tending  to  go  over 
into  the  stable  form  under  the  influence  of  boiling  alcohol 
or  during  the  process  of  a  chemical  reaction.  This  action 
of  alcohol  increases  the  difficulty  of  isolating  the  labile  form 
and  naturally  suggests  the  idea  of  using  other  solvents ;  but 
imperfect  as  was  the  separation  of  the  two  isomers,  better 
results  were  obtained  from  an  alcoholic  solution  than  by 
any  other  means. 

The  phenomena  described  in  the  article  to  which  reference 
was  previously  made,  have  been  reproduced  completely  in  the 
ethers  made  by  the  action  of  acid  chlorides  on  the  sodium  salt 
of  the  oxime  (to  be  described  later)  produced  by  the  action 
of  pure  amyl  nitrite  on  the  sodium  salt  of  the  cresol.  In  the 
resulting  metathesis,  as  hi  the  product  of  the  reaction  of 
hydroxylamine  on  the  quinones,  there  is  not  the  possibility  of 
the  formation  of  a  nitro-body  as  there  is  in  the  action  of  free 
nitrous  acid  on  the  cresols.  Consequently  the  interfering 
action  of  an  admixture  of  such  a  body  with  the  oxime  ethers 


OF  THE   TOLUdUINONEOXIME  ETHERS.  285 

cannot  influence  the  results  as  obtained,  yet  the  presence  of  a 
low-melting  body  existing  hi  a  larger  proportion  in  this  product 
than  in  the  substances  formed  by  the  agency  of  free  nitrous 
acid  according  to  the  type  fraction,  was  indisputably  evidenced. 
The  identity  of  the  observed  phenomena  under  varying 
conditions  and  methods  of  formation,  thus  removing  the 
probability  of  an  admixture  of  an  impurity  and  at  the  same 
tune  of  a  structural  difference,  seems  to  establish  the  hypothesis 
of  a  space  isomerism  in  the  case  of  the  ethers  of  toluquinone- 
metaoxime. 

Inasmuch  as  the  monohalogen  derivatives  of  the  quinone- 
oxime  ethers  are  beautifully  crystalline  bodies,  it  was  thought 
advisable  to  prepare  the  monochlor-  and  monobrombenzoyl 
ethers  of  toluquinoneorthooxime  in  the  hope  of  obtaining 
from  these  well-characterized  products  additional  evidence  as 
to  the  existence  of  isomeric  phenomena  in  the  orthooxime 
ethers.  Although  there  was  the  possibility  of  both  "  space" 
and  "  place  "  isomerism  neither  body  offered  any  indication  of 
the  presence  of  isomers  of  any  kind,  but  each  appeared  to  be 
an  entirely  homogeneous  and  simple  substance,  liquefying 
sharply  at  a  definite  melting-point. 

EXPERIMENTAL  PART. 
The  Sodium  Salt  of  Toluquinone-m-oxime. 

The  sodium  salt  was  prepared  according  to  the  general 
method  suggested  by  Walker.*  To  a  molecule  of  sodium 
alcoholate  freshly  prepared  by  dissolving  metallic  sodium  in 
as  little  alcohol  as  possible,  a  molecule  of  the  orthocresol  was 
added  and  the  solution  treated  with  slightly  more  than  the 
theoretical  quantity  of  amyl  nitrite,  the  whole  being  thoroughly 
mixed  together.  On  standing  in  a  vacuum  over  sulphuric 
acid,  the  sodium  salt  separates  in  fine  purple  crystals,  which, 
on  washing  carefully  with  ether  to  remove  amyl  alcohol  and 
excess  of  nitrite,  is  ready  for  use.  It  may  be  recrystallized 
from  dilute  alcohol  if  further  purification  is  desired.  This 

*  Walker,  Ber.  Dtsch.  chem.  Ges.,  xvii,  399. 


286  NOTES  ON  THE  SPACE  ISOMERISM 

salt  is  extremely  soluble  in  water,  much  less  in  alcohol,  and 
insoluble  in  most  other  organic  liquids.  On  standing  in  the 
air  it  tends  to  decompose,  turning  almost  black.  On  analysis : 

0.1231  gram,  dried  over  H2S04,  gave  0.0541  gram  Na2S04. 

Calculated  for  Fmind 

C7He02NNa.  Found. 

Na  14.48  14.26 


Monobromtoluquinone-o-oxime  Benzoyl  Ether. 

This  ether  was  obtained  from  the  dibromide  previously 
described*  by  boiling  with  dilute  75  per  cent  alcohol,  during 
which  process  hydrobromic  acid  is  split  off.  Being  much  less 
soluble  in  alcohol  than  the  dibromide,  the  monobrom-body 
separates  from  the  solution  as  soon  as  formed  and  may  be 
obtained  as  a  yellow  crystalline  powder  by  filtration  of  the 
cooled  liquid.  On  analysis, 

0.1450  gram,  dried  over  H2S04,  gave  0.2774  gram  C02,  and 
0.0415  gram  H20. 

0.1102  gram  gave  0.0643  gram  AgBr. 

Calculated  for  w        , 

CMH10BrN08. 

C  52.49  52.17 

H  3.15  3.18 

Br  24.98  24.83 

On  crystallizing  from  alcohol  the  ether  was  readily  obtained 
in  two,  apparently  unlike,  modifications,  one  being  long 
prismatic  crystals,  the  other  appearing  as  broad,  thick,  mono- 
clinic  plates.  It  was  at  first  believed  that  this  distinction  in 
crystal  form  was  due  to  the  presence  of  isomeric  bodies,  but 
no  difference  in  melting-point  could  be  found  since  each 
portion  liquefied  sharply  at  184°.  Under  the  lens,  moreover, 
the  plates  are  seen  to  be  striated  parallel  to  one  edge  and  have 
all  the  appearance  of  consisting  of  a  number  of  the  simple 
crystals  united  to  each  other,  since  both  forms  are  plainly  of 

*  Bridge  and  Morgan,  Am.  Chem.  Jour.,  xx,  776. 


OF  THE   TOLUQUINONEOXIME  ETHERS.  287 

the  same  system.     Slow  cooling  was  found  to  be  productive 
of  the  massed  crystals. 

Toluquinone-o-oxime  Benzoyl  Ether  Dichloride. 

The  benzoyl  ether  was  dissolved  in  a  small  amount  of 
chloroform  and  dry  chlorine  gas  passed  into  the  solution. 
The  action  is  very  rapid  and  after  fifteen  minutes  the  liquid 
may  be  allowed  to  evaporate  and  the  white  product,  crystallized 
once  from  glacial  acetic  acid,  melts  sharply  to  a  colorless 
liquid  at  149°  without  decomposition.  Fractional  crystalliza- 
tion does  not  change  the  melting-point,  except  as  the  action  of 
the  solvent  causes  a  slight  formation  of  the  monochlor-body. 
On  analysis, 

0.1241  gram,  dried  over  H2S04,  gave  0.2431  gram  C02,  and 
0.0424  gram  H2O. 

0.2105  gram  gave  0.1930  gram  AgCl. 


Calculated  for  ,,       . 

Found- 


C  53.84  53.42 

H  3.55  3.80 

Cl  22.72  22.67 

The  dichloride,  like  the  dibromide,  is  slightly  soluble  in 
alcohol,  readily  soluble  in  chloroform,  glacial  acetic  acid,  and 
fuming  nitric  acid.  Water  precipitates  the  ether  entirely 
unchanged  from  the  two  last-mentioned  solvents.  In  crystal 
form  it  resembles  very  closely  the  dibromide,  separating  from 
a  boiling  acetic  acid  solution  in  short,  thick,  colorless,  almost 
microscopic  prisms,  suggesting  the  orthorhombic  system.  In 
point  of  stability  this  ether,  as  well  as  the  monochlor-body, 
far  surpasses  the  corresponding  bromine  compounds. 

Monochlortoluquinone-o-oxime  Benzoyl  Ether. 

Analogous  to  the  dibromide,  by  the  action  of  dilute  alcohol 
on  the  dichloride,  hydrochloric  acid  is  split  off  and  a  yellow 
monochlor-substitution-product  is  formed,  one  hour  being 
sufficient  to  complete  the  reaction.  This  ether  closely  resem- 


288      ISOMERISM  OF  TOLUQUINONEOXIME  ETHERS. 

bles  the  monobrom-derivative  in  properties  and  in  crystal 
form.  The  product  as  obtained  from  an  alcoholic, solution 
will  decompose  and  melt  at  a  temperature  varying  ordinarily 
from  185°-193°,  depending  on  the  rapidity  with  which  the 
heat  is  applied.  "  Dipped "  for  fifteen  seconds,  the  crystals 
melt  without  decomposition  at  200°.  Although  carefully 
fractioned  from  alcohol,  ligroin,  and  benzol,  the  substance 
appeared  to  be  entirely  homogeneous,  and  no  variation  in  the 
melting-point  was  observed.  On  analysis, 

0.1209  gram,  dried  over  H2S04,  gave  0.2688  gram  C02,  and 
0.0424  gram  H20. 

0.1035  gram  gave  0.0533  gram  AgCl. 

Calculated  for  -prt,,«j 

CMH10C1N08.  Found< 

C  60.97  60.63 

H  3.66  3.90 

Cl  12.87  12.75 


XXXVI 

ON  THE  VOLUMETRIC   ESTIMATION  OF 
CERIUM. 

BY  PHILIP  E.  BROWNING.* 

SOME  forty  years  ago  Bunsenf  showed  that  the  oxide 
obtained  by  the  ignition  of  cerium  oxalate  might  be  estimated 
volumetrically  by  bringing  it  in  contact  with  potassium  iodide 
and  strong  hydrochloric  acid  and  determining  the  iodine  set 
free.  This  method  may  be  briefly  described  by  a  translation 
of  part  of  the  original  article :  "  The  substance  to  be  deter- 
mined is  weighed  out  in  a  glass  flask  of  from  ten  to  fifteen 
cubic  centimeters  capacity,  a  few  crystals  of  potassium  iodide 
are  added,  and  the  neck  of  the  flask  is  drawn  out  by  the  aid  of 
a  blowpipe  to  a  narrow  opening.  The  flask  is  filled  almost  to 
the  narrowing  of  the  neck  with  hydrochloric  acid  which  is  free 
from  chlorine  or  iron  chloride,  and  a  little  sodium  carbonate  is 
added  in  order  to  displace  the  last  trace  of  air  by  carbon 
dioxide.  The  flask  is  then  closed  by  sealing  off  the  neck  in 
the  blowpipe  and  warmed  in  a  water  bath  until  the  cerium 
compound  is  completely  dissolved,  and  the  quantity  of  iodine 
set  free  is  determined  by  iodometric  analysis." 

The  anhydrous  dioxide  prepared  by  the  ignition  of  the 
oxalate  or  hydroxide  is  very  slowly  acted  on  by  acids, 
especially  when  pure.f  For  this  reason  the  method  which 
Bunsen  described  has  remained  the  only  one  adapted  to  the 
satisfactory  volumetric  estimation  of  the  ignited  dioxide. 

Two  portions  of  the  dioxide  were  prepared  by  treating  the 
crude  cerium  chloride  in  concentrated  solution  with  gaseous 

*  From  Am.  Jour.  Sci.,  viii,  451. 

t  Ann.  Chem.  Phar.,  cv,  49. 

$  Rose,  Handbuch  der  analytischen  Chemie,  Band  i,  219. 

VOL.  II.  — 19 


290  VOLUMETRIC  ESTIMATION  OF  CERIUM. 

hydrochloric  acid  *  to  saturation  to  remove  the  iron.  The 
cerium  chloride  was  then  dissolved  hi  water,  potassium 
hydroxide  added  in  excess  and  chlorine  gas  passed  until  the 
precipitate  became  distinctly  orange  in  color  and  the  solution 
gave  a  strong  odor  of  chlorine.!  This  operation  was  repeated 
until  a  portion  of  the  precipitate  dissolved  in  acid  showed 
no  didymium  absorption  bands  when  examined  before  the 
spectroscope.  The  whole  precipitate  of  the  dioxide  was  then 
dissolved  in  hydrochloric  acid  and  the  oxalate  precipitated  by 
ammonium  oxalate  in  large  excess.  The  precipitated  oxalate 
was  then  washed  thoroughly  with  hot  water  until  the  washings 
gave  no  test  for  hydrochloric  or  oxalic  acids  and  ignited  to  the 
dioxide.  Another  portion  of  the  dioxide  was  later  prepared 
by  precipitating  a  solution  of  pure  cerium  chloride  by  means 
of  ammonium  oxalate,  washing  and  igniting  as  described. 
The  dioxide  in  all  three  cases  was  of  a  light  chamois  color, 
and  uniform  results  were  obtained  from  the  three  portions. 


A  modification  of  the  method  of  Bunsen  —  (with  G.  A. 
FOKD  and  F.  J.  HALL). 

Weighed  portions  of  the  pure  cerium  dioxide  were  placed  in 
small  glass  stoppered  bottles  of  about  100  cm3  capacity,  together 
with  a  gram  of  potassium  iodide  free  from  iodate  and  a  few 
drops  of  water  to  dissolve  the  iodide.  A  current  of  carbon 
dioxide  was  passed  into  the  bottle  for  about  five  minutes  to 
expel  the  air,  10  cm3  of  pure  strong  hydrochloric  acid  were 
added,  the  stopper  inserted  and  the  bottle  heated  gently  upon 
a  steam  radiator  for  about  one  hour  until  the  dioxide  dissolved 
completely  and  the  iodine  was  set  free.  After  cooling  the 
bottle,  to  prevent  loss  of  iodine  upon  removing  the  stopper, 
the  contents  were  carefully  washed  into  about  400  cm3  of 
water  and  titrated  with  standard  sodium  thiosulphate  to 
determine  the  amount  of  iodine  liberated  according  to  the  well 
known  reaction 

2Ce02  +  8HC1  +  2KI  =  2CeCls  +  2KC1  +  4H20  +  I2. 

*  Dennis  and  Magee,  Zeitschr.  anorg.  Chem.,  iii,  260. 

t  Mosander,  Phil.  Mag.,  xxviii,  241  ;  Dennis,  Zeitschr.  anorg.  Chem.,  yii,  252. 


VOLUMETRIC  ESTIMATION  OF  CERIUM. 


291 


A  few  blank  determinations  were  carried  through  in  the  bottles 
without  the  presence  of  the  cerium  dioxide  to  determine  the 
amount  of  iodine  set  free  under  these  conditions.  The  amount 
obtained  was  uniformly  equal  to  0.04  cm3  of  the  ^  iodine 
solution  which  was  taken  as  the  correction  and  applied  to  all 
the  determinations.  The  results  follow  in  Table  I. 

TABLE  I. 


Exp. 

CeO,  taken. 

CeOj  found. 

Error. 

grm. 

grm. 

grm. 

(1) 

0.1000 

0.0994 

0.0006- 

(2) 

0.1032 

0.1034 

0.0002+ 

(3) 

0.1016 

0.1017 

0.0001+ 

(4) 

0.1054 

0.1041 

0.0013- 

(5) 

0.2010 

0.2021 

0.0011+ 

(6) 

0.1104 

0.1109 

0,0005+ 

(7) 

0.1914 

0.1907 

0.0007- 

(8) 

0.1604 

0.1603 

0.0001- 

(9) 

0.2146 

0.2145 

0.0001- 

(10) 

0.1108 

0.1099 

0.0009- 

(11) 

0.1346 

0.1347 

0.0001+ 

(12) 

0.1540 

0.1534 

0.0006- 

13 

0.1976 

0.1968 

0.0008- 

(14) 

0.1230 

0.1240 

0.0010+ 

(15) 

0.1199 

0.1202 

0.0003+ 

(16) 

0.1524 

0.1528 

0.0004+ 

(17) 

0.1212 

0.1211 

0.0001- 

(18) 

0.1528 

0.1543 

0.0015+ 

In  order  to  obtain  a  further  check  upon  the  accuracy  of  the 
method,  portions  of  the  cerium  dioxide  were  weighed  out  and 
placed  in  a  distillation  apparatus  previously  employed  for 
similar  purposes  and  described  in  former  articles  from  this 
laboratory,  viz.:  a  Voit  flask,  serving  as  a  retort,  sealed  to 
the  inlet  tube  of  a  Drexel  wash-bottle,  used  as  a  receiver, 
the  outlet  tube  of  which  was  trapped  by  sealing  on  Will  and 
Varrentrapp  absorption  bulbs.  In  the  retort  the  cerium 
dioxide  together  with  15  cm3  of  water,  1  grm.  of  potassium 
iodide  and  10  cm3  of  pure  strong  hydrochloric  acid  were  placed. 
In  the  receiver  were  100  cm3  of  water  and  2  to  3  grm.  of 
potassium  iodide,  and  in  the  bulbs  a  dilute  solution  of  potas- 
sium iodide.  Before  adding  the  hydrochloric  acid  a  current  of 
carbon  dioxide  was  passed  through  the  apparatus  for  some 


292 


VOLUMETRIC  ESTIMATION  OF  CERIUM. 


minutes.  After  adding  the  acid,  the  liquid  was  boiled  in  the 
current  of  carbon  dioxide  *  to  a  volume  of  15  cm3,  when  the 
free  iodine  had  almost  completely  left  the  retort  and  passed 
into  the  receiver,  and  the  apparatus  was  allowed  to  cool. 

The  iodine  in  the  receiver  was  titrated  directly  with  sodium 
thiosulphate,  and  that  in  the  retort  after  dilution  of  the 
residue  to  about  400  cm3,  the  later  amount  seldom  exceeding 
the  equivalent  of  a  few  drops  of  -^  iodine  solution. 

The  results  f  oUow  in  Table  II. 

Here  also  blank  determinations  were  made  but  no  correction 
was  found  to  be  necessary. 

An  attempt  early  in  the  work  to  titrate  by  an  alkaline  arse- 
nite  the  iodine  liberated,  after  neutralizing  the  hydrochloric 
acid,  brought  out  some  curious  results  which  seem  worthy  of 
mention. 

TABLE  H. 


Exp. 

CeO,  taken. 

CeOj  found. 

Error. 

grm. 

grm. 

grm. 

1) 

0.1028 

0.1013 

0.0015- 

2) 

0.2060 

0.2055 

0.0005- 

3) 

0.2014 

0.2012 

0.0002- 

4 

0.1716 

0.1711 

0.0005- 

5) 

0.0974 

0.0972 

0.0002- 

(6) 

0.1600 

0.1587 

0.0013- 

(7) 

0.1268 

0.1254 

0.0014- 

. 

8) 

0.1276 

0.1268 

0.0008- 

9 

0.1620 

0.1612 

0.0008- 

(10) 

0.1016 

0.1011 

0.0005- 

(H) 

0.1648 

0.1543 

0.0005- 

(12) 

0.1352 

0.1342 

0.0010- 

In  these  experiments  the  contents  of  the  bottles  after  the 
cerium  had  dissolved  were  carefully  washed  into  a  Drexel 
wash  bottle  upon  the  inlet  tube  of  which  was  fused  a  thistle 
tube  with  a  stop-cock  and  to  the  outlet  tube  a  Will  and  Var- 
rentrapp  absorption  trap.  In  the  trap  a  solution  of  potassium 
iodide  was  placed  and  through  the  thistle  tube  a  saturated 

*  The  carbon  dioxide  gas  was  furnished  by  a  Kipp  generator  from  marble 
and  hydrochloric  acid  of  one-half  strength,  both  of  which  had  been  boiled 
previously  to  remove  all  air. 


VOLUMETRIC  ESTIMATION  OF  CERIUM. 


293 


solution  of  potassium  bicarbonate  was  added  to  complete  neu- 
tralization of  the  acid.  Any  iodine  carried  mechanically  by 
the  carbon  dioxide  should  be  held  by  the  potassium  iodide 
solution  in  the  trap.  After  neutralization  the  free  iodine  was 
titrated  by  standard  arsenious  oxide  solution.  The  results 
appear  in  Table  III. 

TABLE  IIL 


Exp. 

CeO2  taken. 

CeO,  found. 

Error. 

grin. 

gnn. 

gnn. 

(1) 

0.1000 

0.0987 

0.0013-] 

2 
(3) 

0.1005 
0.1030 

0.0981 
0.1009 

0.0024-1  T 
0.0021-  [*• 

(4) 

0.1500 

0.1475 

0.0025-  1 

(6) 

0.1030 

0.1005 

0.0025-' 

(6 

0.1010 

0.0988 

0.0022- 

7 

0.1510 

0.1608 

0.0002- 

TT 

(8) 

0.1530 

0.1485 

0.0045- 

•  11. 

(9) 

0.2045 

0.2011 

0.0034- 

(10) 

0.2000 

0.1958 

0.0042- 

(11) 

0.1334 

0.1302 

0.0032- 

(12) 

0.1354 

0.1330 

0.0024- 

(13) 

0.1312 

0.1294 

0.0018- 

(14) 
(15) 

0.1308 
0.1060 

0.1277 
0.1042 

0.0031- 
0.0018- 

m. 

(16) 

0.1602 

0.1567 

0.0035- 

(17) 

0.1504 

0.1488 

0.0016- 

As  will  be  seen  by  the  table,  an  average  error  of  about  2  per 
cent  runs  through  the  entire  set.  The  natural  conclusion 
would  be  that  the  cerium  dioxide  contained  some  impurity; 
but,  as  the  first,  second,  and  third  samples,  very  carefully  pre- 
pared, gave  the  same  results,  it  seemed  necessary  to  look  else- 
where for  an  explanation.  Two  possible  causes  suggested 
themselves :  first,  mechanical  loss  during  the  process  of  neu- 
tralization, and  second,  the  possible  formation,  under  the  condi- 
tions, of  iodine  chloride,  which  if  formed  would  in  the  process 
of  neutralization  probably  take  the  form  of  potassium  chloride, 
iodide,  and  iodate,  and  thus  some  of  the  originally  free  iodine 
would  be  withdrawn  from  the  amount  titrated.  To  test  these 
theories,  portions  of  the  ^  iodine  solution  roughly  equiva- 
lent to  the  amounts  of  iodine  set  free  by  0.1  and  0.2  grms.  of 
CeOa,  were  drawn  off  into  bottles  previously  rilled  with  carbon 


294 


VOLUMETRIC  ESTIMATION  OF  CERIUM. 


dioxide,  treated  with  the  usual  amount  of  strong  hydrochloric 
acid  (10  cm3),  and  after  standing  from  thirty  to  forty-five 
minutes,  neutralized  and  titrated  as  already  described.  The 
results  were  most  interesting  and  seemed  to  show  a  loss  of 
iodine  closely  equivalent  to  that  shown  by  the  results  of 
Table  III,  and  proportional  to  the  amount  of  iodine  originally 
present.  A  few  determinations  were  carried  through  in  the 
same  way  except  that  the  neutralization  was  omitted  and 
dilution  and  titration  with  thiosulphate  substituted.  These 
showed  a  loss  of  iodine  well  within  the  limits  of  such  a  proc- 
ess. The  results  follow  in  Table  IV. 

TABLE  IV. 
WITH  ARSENIOUS  OXIDE. 


Exp. 

Iodine  *>. 
taken. 

Iodine  ,£ 
found. 

Error. 

Equivalent  error 
on  Ce02. 

(1) 
(2) 
(3) 
(4) 
5) 
(6) 
(7) 

cm3 

5.22 
6.09 
5.10 
6.66 
5.10 
10.22 
10.21 

cms 
6.07 

4.97 
4.97 
6.47 
4.97 
9.97 
9.97 

cm8 
0.15- 
0.12- 
0.13- 
0.19- 
0.13- 
0.25- 
0.24- 

grm. 

0.0026- 
0.0021- 
0.0022- 
0.0033- 
0.0022- 
0.0043- 
0.0041- 

WITH  SODIUM  THIOSULPHATE. 

g| 

i 

5 
5 
10 
10.08 

5.01 
4.99 
10.02 
10.12 

0.01+ 
0.01- 
0.02+ 
0.04+ 

0.0002+ 
0.0002- 
0.0003+ 
0.0007+ 

The  action  of  arsenious  oxide  upon  Cerium  Dioxide  —  (with 
WM.  D.  GUTTER). 

The  fact  that  cerium  dioxide  is  reduced  by  hydriodic  acid 
suggested  the  possibility  of  the  application  of  arsenious  acid 
in  acid  solution  to  the  same  end  according  to  the  reaction 

4Ce02  +  As208  =  2Ce203  -f  As206. 

The  extreme  difficulty  with  which  the  ignited  cerium  dioxide 
when  pure  dissolves  in  acids  has  already  been  mentioned,  and 
for  this  reason  it  was  found  practically  impossible  to  obtain 


VOLUMETRIC  ESTIMATION  OF  CERIUM. 


295 


any  results  by  this  method.  Weighed  portions  of  the  dioxide 
were  placed  in  Erlenmeyer  beakers  with  an  excess  of  a  solu- 
tion of  arsenious  oxide  £,  10  cm3  of  (1 : 1)  sulphuric  acid 
were  added,  and  the  boiling  continued  until  the  fuming  point 
of  the  acid  was  reached ;  but  even  at  this  point  only  a  partial 
solution  of  the  dioxide  had  taken  place. 

The  dark  brown  powder  obtained  by  igniting  the  carefully 
washed  oxalates,  precipitated  in  acid  solution  by  treating  a 
solution  of  crude  cerium  chloride  with  ammonium  oxalate  or 
oxalate  acid  is  very  fairly  soluble  in  acids.  Mengel*  has 
recently  shown  that  this  product  contains  a  dioxide  of  praseo- 
didymium  which  acts  as  does  cerium  dioxide  toward  reducing 
agents.  This  fact  makes  the  results  recorded  in  the  treatment 
of  this  ignited  mixture  of  oxides  of  no  value  analytically,  but 
of  interest  in  the  comparative  study  of  the  two  reducing 
agents,  arsenious  oxide  and  hydriodic  acid.  Two  portions  of 
this  mixture  of  oxides  gave  the  following  results,  which  agree 
fairly  well  with  those  of  Mengel. 


Ezp. 

Amount  of 
substance  taken. 

Ce02f-HPr02?) 

Calculated  on 
0.1000  grm. 

8 

grm. 
0.1037 
0.1034 

grm. 

0.0530 
0.0538 

gnu. 
0.0511 
0.0520 

The  average  of  these  results  was  taken  as  a  standard 
—0.0515  grm.  CeO2,  etc.,  to  every  0.1000  grm.  of  material. 
Three  carefully  weighed  portions  of  this  same  material  were 
placed  in  Erlenmeyer  beakers  with  10  cm3  of  £  arsenious 
oxide  solution  and  10  cm3  of  dilute  (1 :  4)  sulphuric  acid  and 
boiled  until  complete  solution  had  taken  place.  The  liquid 
was  then  cooled,  neutralized  with  potassium  bicarbonate  and 
titrated  with  standardized  iodine  to  determine  the  amount  of 
arsenious  oxide  remaining,  and  from  it  the  amount  used  in  the 
reduction  of  the  dioxide  according  to  the  reaction  given  above. 
The  results  obtained  follow. 


*  Zeitschr.  anorg.  Chem.,  xix,  67. 


296 


VOLUMETRIC  ESTIMATION  OF  CERIUM. 


Bxp. 

Amount) 
taken. 

Amount  CeOs 

found. 

CeO,  calculated 
for  0.1000  grm. 

gnu. 

grin. 

grm. 

(1) 

0.1005 

0.0493 

0.0491 

(2) 
(3) 

0.1016 
0.1005 

0.0494 
0.0486 

0.0487 
0.0484 

As  will  be  seen,  the  results  obtained  by  this  method  fall 
about  0.0030  grm.  below  the  standard  as  obtained  by  the  dis- 
tillation method,  which  seems  to  show  that  the  arsenious  oxide 
does  not  effect  the  complete  reduction  of  the  cerium  dioxide 
from  CeO2  to  Ce2O8. 

In  order  to  study  this  point  a  little  more  fully  and  upon  the 
pure  dioxide,  definite  portions  of  a  standard  solution  of  pure 
cerium  chloride  were  precipitated  by  ammonia  in  the  presence 
of  hydrogen  dioxide  and  boiled  to  reduce  the  CeO8  formed  to 
the  conditions  of  CeO2.  The  precipitated  hydrated  dioxide 
was  filtered  off  and  carefully  washed  until  the  washings  gave 
no  indication  of  hydrogen  dioxide.  The  moist  precipitate  was 
then  washed  into  a  beaker,  one  gram  of  potassium  iodide  added 
and  10  cm3  of  strong  HC1.  The  precipitate  dissolved  quite 
readily  in  the  cold  and  the  iodine  liberated  was  determined  by 
standard  sodium  thiosulphate.  The  results  appear  in  Table  V. 

TABLE  V. 


Bxp. 

CeO2  taken. 

CeO2  found. 

Error. 

grm. 

grm. 

grm. 

1) 

0.1142 

0.1140 

0.0002- 

2) 

0.1142 

0.1147 

0.0005+ 

i 

0.1142 
0.1142 

0.1152 
0.1159 

0.0010+ 
0.0017+ 

(5) 

0.1142 

0.1152 

0.0010+ 

0.1142 

0.1156 

0.0014+ 

Another  series  of  these  precipitates  prepared  in  the  same 
way  was  boiled  with  a  definite  amount  of  arsenious  acid  in 
acid  solution,  as  previously  described  in  the  case  of  the 
ignited  dioxide.  The  results  which  are  recorded  in  Table  VI 


VOLUMETRIC  ESTIMATION  OF  CERIUM. 


297 


show,  as  in  the  case  of  the  ignited  dioxide,  an  insufficient 
reduction  of  the  cerium  by  the  arsenious  acid. 


TABLE  VI. 


Exp. 

CeO2  taken. 

Ce08  found. 

Error. 

gnn. 

grin. 

grin. 

(1) 

0.0881 

0.0370 

0.0011- 

2) 

0.0381 

0.0361 

0.0020- 

(3) 

0.1142 

0.1077 

0.0064- 

(4) 

0.1060 

0.1002 

0.0058- 

The  Estimation  of   Cerium    Oxalate  ly  Potassium  Perman- 
ganate—  (with  LEO  A.  LYNCH). 

Stolba*  has  stated  that  cerium  oxalate  may  be  estimated 
volumetrically  after  the  same  manner  as  calcium  oxalate  by 
treating  the  washed  precipitate,  suspended  in  warm  water,  to 
which  a  moderate  amount  of  sulphuric  acid  has  been  added, 
by  potassium  permanganate.  As  the  titration  proceeds  the 
precipitate  disappears  and  the  end  reaction  is  sharp.  He  also 
finds  that  the  permanganate  does  not  oxidize  the  cerium 
from  the  lower  to  the  higher  condition.  So  far  as  we  have 
been  able  to  discover,  no  experimental  evidence  has  been 
presented  to  prove  the  correctness  of  Stolba's  statement,  and 
the  work  to  be  described  was  undertaken  to  furnish  such 
evidence. 

The  solutions  used  were  prepared  and  standardized  as 
follows :  The  cerium  solutions  were  made  by  dissolving 
10  grams  of  pure  cerium  chloride  in  one  liter  of  water,  and 
standardized  by  precipitating  measured  and  weighed  portions, 
in  a  faintly  acid  solution,  with  ammonium  oxalate,  filtering, 
washing,  igniting,  and  weighing  as  the  dioxide  (CeO2).  A 
solution  of  potassium  permanganate  was  prepared  and  stan- 
dardized by  titration  against  weighed  amounts  of  ammonium 
oxalate.  A  solution  of  ammonium  oxalate  was  made  and 
its  value  determined  by  titrating  measured  amounts  against 

*  Sitzungsber.  d.  kgl.  bohm.  Gesellsck.  d.  Wissenschaften  v.  4.  Juli,  1879; 
Zeitschr.  anal.  Chem.,  xix,  194. 


298          VOLUMETRIC  ESTIMATION  OF  CERIUM. 

potassium  permanganate.  Definite  portions  of  the  cerium 
solution  were  drawn  from  a  burette  and  after  diluting  with 
water  from  100  to  200  cm3  a  definite  amount  of  ammonium 
oxalate  was  added,  care  being  taken  to  have  an  excess  over 
the  amount  necessary,  and  the  whole  warmed  to  insure  a 
more  crystalline  precipitation.*  The  precipitate  was  then 
filtered  off  on  paper  and  carefully  washed,  the  filtrate  and 
washings  being  collected  in  a  liter  Erlenmeyer  flask  and  set 
aside  for  future  use.  The  precipitate  was  treated  with  about 
10  cm3  of  hot  (1 :  4)  sulphuric  acid,  which  dissolved  it 
completely,  if  not  at  first,  by  running  it  through  the  filter 
a  few  times,  and  the  solution  and  washings  were  collected  hi 
another  liter  flask.  The  total  volume  of  liquid  was  made 
up  to  about  500  cm3,  warmed  to  about  70°  C.  to  80°  C.  and 
titrated  with  potassium  permanganate  to  the  appearance  of 
the  faint  blush  of  color  showing  the  complete  oxidation  of  the 
oxalic  acid.  The  filtrate  from  the  cerium  oxalate  containing 
the  excess  of  oxalic  acid  was  diluted  to  500  cm3,  acidified 
with  10  cm3  of  dilute  (1  :  4)  sulphuric  acid,  one  gram  of 
manganous  sulphate  added  to  prevent  the  interfering  action 
of  the  free  hydrochloric  acid  upon  the  estimation  of  the  oxalic 
acid,f  and  titrated  with  potassium  permanganate  after  the 
same  manner  as  the  dissolved  precipitate.  A  definite  quantity 
of  ammonium  oxalate  having  been  originally  taken,  it  became 
possible,  by  subtracting  from  it  the  amount  obtained,  to 
derive  the  measure  of  the  oxalate  used  in  the  precipitation 
of  the  cerium  oxalate.  By  this  procedure,  it  will  be  observed 
a  check  was  made  upon  the  results  obtained  by  the  titration 
of  the  precipitate.  In  experiments  (1)  to  (6)  the  cerium 
oxalate  was  thrown  down  in  neutral  solution,  in  experiments 
(7)  to  (10)  in  acid  solutions.  The  treatment  of  the  filtrate 
in  experiment  (1)  was  made  without  the  presence  of  the 
manganous  sulphate.  The  results  recorded  in  Table  VII 
seem  to  uphold  the  statement  of  Stolba. 

*  As  shown  by  the  table,  the  precipitation  was  sometimes  in  neutral,  some- 
times in  faintly  acid  solution. 

t  Gooch  and  Peters,  Am.  Jour.  Sci.,  vii,  461.    This  volume,  p.  222. 


VOLUMETRIC  ESTIMATION  OF  CERIUM.          299 
TABLE  VII. 


Bxp. 

Amount  found. 

Amount  found. 

Amount  taken. 

Calculated  as 

Error. 

Calculated  as 

Error. 

Calculated 

CeCl3. 

Calculated  as 

CeCl8. 

Calculated  as 

as  CeCl3. 

Treatment  of 

CeCl3. 

Treatment  of 

CeCl8. 

precipitate. 

filtrate. 

grm. 

grm. 

grlu. 

grm. 

grm. 

(1 

[0.1091 

0.1087 

0.0004- 

0.1023 

0.0068-] 

(2 

0.1091 

0.1103 

0.0012+ 

. 

g 

3 

0.1091 

0.1087 

0.0004- 

0.1087 

0.0004- 

4 

0.1364 

0.1373 

0.0009+ 

0.1391 

0.0027+ 

1: 

0.1364 
0.2182 

0.1367 
0.2202 

0.0003+ 
0.0020+ 

0.1367 
0.2206 

0.0003+ 
0.0024+ 

(7) 

0.1091 

0.1087 

0.0004- 

(8) 

0.1519 

o.is35 

0.0016+ 

0.1535 

0.0016+ 

(9) 

0.1364 

0.1367 

0.0003+ 

0.1367 

0.0003+ 

(10) 

0.2182 

0.2183 

0.0001+ 

0.2183 

0.0000 

XXXVII 

ON    THE    ESTIMATION   OF    THALLIUM    AS  THE 
CHROMATE. 

BY  PHILIP  E.  BROWNING  AND  GEORGE  P.  HUTCHINS.* 

CROOKES  has  shown  f  that  the  chromate  precipitated  by  the 
addition  of  potassium  dichromate  to  an  alkaline  solution  of  a 
thallous  salt  has  the  constitution  of  a  neutral  salt  and  is  very 
insoluble  in  water  — 100  parts  of  water  at  100°  C.  dissolving 
about  0.2  parts  and  at  60°  C.  about  0.03  parts.  He  has  also 
made  use  of  this  reaction  J  to  effect  a  rough  separation  of 
thallium  from  cadmium. 

The  object  of  this  paper  is  to  describe  some  work  directed 
toward  a  study  of  the  application  of  this  reaction  to  the  gravi- 
metric estimation  of  thallium  and  the  best  conditions  under 
which  to  effect  the  precipitation.  For  the  work  a  solution 
of  thallous  nitrate  was  made  by  dissolving  10  grms.  in  water 
and  making  up  to  a  liter.  The  standard  was  determined 
by  taking  measured  and  weighed  portions  from  a  burette, 
precipitating  with  a  slight  excess  of  potassium  iodide,  agitat- 
ing to  bring  about  a  good  separation  of  the  thallous  iodide, 
and  allowing  to  stand  until  the  supernatant  liquid  was 
clear.  The  iodide  was  then  filtered  off  upon  an  asbestos  felt 
contained  in  a  perforated  platinum  crucible,  the  whole  having 
been  previously  ignited  and  weighed,  washed  with  a  mixture 
of  alcohol  and  water,  dried  over  a  low  flame  and  weighed 
to  a  constant  weight.  The  filtrate,  which  together  with 
the  washings  seldom  amounted  to  more  than  50  cm3,  was 
evaporated  to  dryness  on  a  water  bath,  a  few  drops  of  water 

*  From  Am.  Jour.  Sci.,  viii,  460.  t  Chem.  News,  viii,  255. 

J  Chem.  News,  vii,  145. 


ESTIMATION  OF  THALLIUM  AS   CHEOMATE.       301 

added  and  thus  the  small  amount  of  thallous  iodide  which 
had  been  dissolved  recovered.  This  small  insoluble  residue, 
which  seldom  amounted  to  one  milligram  in  weight,  was 
filtered  off,  washed  and  weighed  as  previously  described. 
Baubigny  *  has  shown  this  method  to  give  very  satisfactory 
results,  and  the  uniformity  of  our  determinations  certainly 
confirms  his  statements. 

For  convenience  in  the  calculations  of  results  to  be  de- 
scribed later,  a  solution  of  potassium  dichromate  of  definite 
strength  was  made.  Portions  of  the  thallium  solution  were 
drawn  from  a  burette  into  test  tubes  of  about  100  cm3  capacity 
and  weighed  as  a  check  on  the  burette  reading.  The  solution 
was  heated  to  about  70°  C.  to  80°  C.  and  a  few  drops  of 
ammonia  or  potassium  carbonate  solution  added  to  distinct 
alkalinity.  A  definite  amount  of  the  potassium  dichromate 
in  solution  was  delivered  from  a  burette,  care  being  taken  to 
have  an  excess,  and  the  contents  agitated  to  bring  about  a 
good  separation  of  the  precipitated  chromate.  After  the 
precipitate  had  completely  settled  out  and  the  solution  had 
become  cold  the  chromate  was  filtered  upon  asbestos,  as 
described  above,  dried  over  a  low  flame  and  weighed  to  a 
constant  weight.  The  filtrates  from  several  determinations 
were  evaporated  to  a  small  volume  and  in  one  or  two  cases 
a  residue  amounting  to  a  few  tenths  of  a  milligram  was 
obtained,  but  no  appreciable  quantity  of  dissolved  chromate 
was  thus  recovered.  It  was  found  that  when  the  precipita- 
tion was  made  in  the  cold  the  chromate  did  not  flock  well, 
but  remained  partly  in  a  finely  divided  condition  which  would 
run  through  the  felt  and  require  repeated  filtration.  The 
addition  of  ammonium  nitrate  before  precipitation  prevented 
this  largely,  even  in  the  cold,  bat  the  best  results  were 
obtained  by  warming  the  solution  before  precipitation  and 
using  potassium  carbonate  rather  than  ammonium  hydroxide. 
The  results  follow  in  Table  I. 

An  attempt  was  made  to  estimate  the  thallium  volumetri- 
cally  by  determining  the  amount  of  chromate  in  the  filtrate 
*  Chem.  News,  Ixiv,  239. 


302      ESTIMATION  OF  THALLIUM  AS  CHROMATE. 

TABLE  I. 


Brp. 

TINOs  taken. 
Calculated  as 
11,0. 

TliCrO4  found. 
Calculated  as 
TlaO. 

Error. 
Calculated  as 
T120. 

grm. 

grm. 

grm. 

(1) 

0.0796 

0.0791 

0.0005- 

(2) 

0.0792 

0.0788 

0.0004- 

(3) 
(4) 

0.0792 
0.1188 

0.0786 
0.1177 

0.0006- 
0.0011- 

(5) 

0.1192 

0.1186 

0.0006- 

(6) 

0.1185 

0.1178 

0.0007- 

(7) 

0.1190 

0.1185 

0.0005- 

(8) 

0.1189 

0.1183 

0.0006- 

(9) 

0.1196 

0.2000 

0.0004+ 

(10) 

0.1196 

0.2005 

0.0009+ 

(ID 

0.1173 

0.1173 

0.0000 

(12) 

0.1171 

0.1163 

0.0008- 

from  the  thallous  chromate,  and  by  difference  (the  potassium 
dichromate  originally  added  being  known)  the  amount 
combined  with  the  thallium  in  the  precipitate.  The  method 
used  to  determine  the  standard  of  the  dichromate  solutions 
and  also  the  chromate  remaining  in  the  filtrate  was  described 
by  one  of  us  in  a  previous  paper  from  this  laboratory.* 
According  to  this  procedure  the  filtrate  from  the  thallous 
chromate  containing  the  excess  of  alkali  chromate  was  acidified 
with  sulphuric  acid,  a  definite  amount  of  a  solution  of  arsenious 
oxide,  previously  standardized,  was  added  and  the  whole  was 
allowed  to  stand  a  few  moments  until  the  change  from  the 
yellow  to  the  bluish  green  showed  the  complete  reduction  of 
the  chromic  acid.  Potassium  bicarbonate  was  added  to  distinct 
alkaline  reaction  and  the  arsenious  oxide  remaining  was 
determined  by  titration  with  -standard  iodine  solution.  The 
amount  of  the  arsenious  oxide  oxidized  is  of  course  the 
measure  of  the  chromate  in  the  solution.  The  amount  of 
chromate  in  the  original  solution  used  being  known,  by 
subtracting  the  amount  thus  determined  in  the  filtrate  the 
chromate  in  combination  with  the  thallium  may  be  readily 
found,  and  from  it  the  thallium  estimated.  Filtrates  from 
certain  precipitates,  of  which  the  determinations  are  given  in 


*  Am.  Jour.  Sci.,  i,  35, 1896.    Volume  I,  p.  344. 


ESTIMATION  OF  THALLIUM  AS  CHROMATE.       303 


Table  I,  were  treated  in  this  way,  and  the  results,  indicated  by 
corresponding  numbers,  follow  in  Table  II. 


TABLE  II. 


Bxp. 

TINO3  taken. 
Calculated  as 

Tl8CrO4  found. 
Calculated  as 

Error. 
Calculated  as 

TlaO. 

Tl,0. 

TljO. 

grm. 

grm. 

grm. 

(5) 

0.1192 

0.1198 

0.0006+ 

(8) 

0.1189 

0.1205 

0.0016+ 

(9) 

0.1196 

0.1180 

0.0016- 

(10) 

0.1196 

0.1192 

0.0004- 

(11) 

0.1173 

0.1182 

0.0009+ 

(12) 

0.1171 

0.1190 

0.0019+ 

The  method  cannot  be  very  accurate  on  account  of  the  high 
molecular  weight  of  thallium  oxide  as  compared  with  that  of 
the  chromic  acid  determined,  but  the  results  check  fairly  well 
with  the  gravimetric  method. 


xxxvm 

THE  ETHERS   OP  ISONITROSOGUAIACOL  IN 

THEIR  RELATION  TO  THE  SPACE 

ISOMERISM  OF  NITROGEN. 

BY  JOHN  L.  BRIDGE  AND  WM.  CONGEE  MORGAN* 

WHEN  the  presence  of  isomerism  in  the  quinoneoximes  was  first 
noted  by  one  of  us,f  and  the  phenomenon  exhibited  by  the 
two  isomeric  ethers  then  described  was  shown  by  Kehrmann  J 
to  be  due  to  no  structural  differences,  but  to  necessitate  the 
assumption  of  a  spatial  arrangement  about  the  nitrogen  atom, 
according  to  the  theory  of  Hantzsch  and  Werner,  the  plan  was 
adopted  of  studying  the  various  substituted  quinoneoximes, 
by  means  of  their  acyl  and  alkyl  ethers,  with  reference  to 
this  phenomenon.  Accordingly  we  have  investigated  the 
toluquinoneoximes,§  both  ortho  and  meta,  producing  them 
by  the  action  of  nitrous  acid  on  the  cresol  as  well  as  by 
hydroxylamine  on  toluquinone,  and  found  that,  whereas  there 
is  abundant  evidence  for  the  existence  of  stereoisomeric  bodies 
in  the  metaoxime  ethers,  in  the  derivatives  of  the  orthooxime, 
all  such  indication  is  wanting.  The  significance  of  this 
observation  is  furthermore  increased  by  the  fact  that  all 
oximes,  in  which  isomerism  has  been  reported,  may  be 
considered  as  derivatives  of  metasubstituted  quinones,  and  it 
seemed  not  improbable,  therefore,  that  these  observations 
might  be  formulated  into  a  general  rule  regarding  the 
appearance  of  isomerism  in  the  quinoneoximes. 

In  the  course  of  his  investigation  of  the  properties  and 
reactions  of  isonitrosoguaiacol,  among  other  derivatives,  Pfob  || 

*  From  Am.  Chem.  Jour.,  xxii,  485.     t  Ann.  Chem.  (Liebig),  cclxxvii,  79. 

t  Ann.  Chem.  (Liebig),  cclxxix,  27. 

§  Am.  Chem.  Jour.,  xx,  761 ;  xxii,  402.    This  volume,  pp.  145,  283. 

II  Monatsh.  Chem.,  xyiii,  467. 


THE  ETHERS  OF  ISONITROSOGUAIACOL  305 

made  the  methyl  and  acetyl  ethers,  but  did  not  announce  the 
observation  of  any  cases  of  isomerism.  Isonitrosoguaiacol  may 
be  considered  as  the  metamethoxyquinoneoxime,  hence  this 
metasubstituted  quinoneoxime  presented  conditions  differing 
widely  from  the  other  members  of  the  same  series  above 
mentioned.  Moreover,  because  of  its  close  relationship  to 
toluquinonemetaoxime,  from  which  it  differs  only  by  the 
interposition  of  an  oxygen  atom  between  the  ring  and  the 
methyl  group,  it  seemed  possible  that  isomeric  modifications 
of  the  ethers  might  exist,  which  had  been  overlooked  by  the 
former  investigator.  When,  furthermore,  Rupe  *  made  no 
mention  of  such  appearance  in  his  research  on  isonitrosoguaia- 
col,  it  seemed  advisable  to  undertake  anew  the  investigation  of 
this  body  with  the  special  purpose  of  discovering  such 
isomerism,  if  possible,  and  to  couple  with  it  an  investigation 
of  the  orthomethoxyquinoneoxune  or  isonitroso  derivative  of 
the  monomethyl  ether  of  resorcin,  in  order  by  this  means  to 
be  able  to  parallel  in  these  closely  analogous  bodies  the 
experiments  with  the  ortho-  and  metacresols. 

With  this  idea,  the  work  of  Pfob  and  Rupe  was  carefully 
repeated,  so  far  as  it  pertained  to  the  question  in  hand, 
but,  aside  from  minor  differences,  our  results  served  only  to 
corroborate  the  testimony  of  these  investigators.  New 
derivatives,  to  be  described  later,  were  prepared  in  the  hope 
that  these  bodies  might  show  some  variations  leading  to  the 
discovery  of  isomeric  modifications,  but  each  appeared  to  be 
entirely  homogeneous  and  no  evidence  for  isomerism  could  be 
found.  These  results  are  of  course  only  negative  and  do  not 
disprove  the  existence  of  space  isomers  in  the  same  bodies, 
yet  the  same  methods,  which  gave  very  positive  evidence  of 
their  presence  hi  other  quinoneoxime  ethers,  were  used  to 
detect  them  in  this  instance. 

Aside  from  the  difficulty  of  obtaining  the  pure  monomethyl 
ether  of  resorcin  in  any  considerable  quantity,  because  of  poor 
synthetical  processes  and  inefficient  methods  of  separation, 
inasmuch  as  preliminary  experiments  pointed  to  a  multiplicity 

*  Ber.  Dtsch.  chem.  Ges.,  xxx,  2444. 
VOL.  ii. —  20 


306          THE  ETHERS  OF  ISONITROSOGUAIACOL. 

of  products  in  the  reaction  with  nitrous  acid  such  as  Kietaibl  * 
found  with  the  monoethyl  ether,  it  was  thought  inadvisable  in 
the  light  of  the  results  obtained  with  the  meta  body,  to  con- 
tinue the  work  on  the  ethers  of  orthomethoxyquinoneoxime. 
Work  along  the  general  line  will  be  continued,  and  the  results 
of  experimentation  with  mononitrosoresorcin  will  soon  appear. 

EXPERIMENTAL  PART. 
Isonitrosoguaiaeol  and  Salts. 

The  isonitrosoguaiacol  used  in  the  investigation  was  pre- 
pared both  by  the  method  of  Pfob,  working  with  nascent  nitrous 
acid  in  alcoholic  solution,  and  also  by  the  general  method  for 
the  formation  of  the  sodium  salts  of  isonitroso  bodies  sug- 
gested by  Walker,  f  To  a  concentrated  alcoholic  solution  of 
sodium  alcoholate,  guaiacol  is  added  in  sufficient  quantity  to 
form  the  sodium  salt  by  the  resulting  metathesis,  then,  to  this 
solution  of  sodium  guaiacol,  slightly  more  than  the  theoretical 
quantity  of  amyl  nitrite  is  added  and,  after  thorough  mixing, 
the  liquid  is  allowed  to  stand  over  sulphuric  acid  for  twenty- 
four  hours,  when  the  bright  olive-green  crystalline  sodium  salt 
of  isonitrosoguaiacol  separates.  After  washing  thoroughly 
with  ether,  pulverizing  and  rewashing,  the  salt  may  be  used 
directly  for  the  preparation  of  derivatives,  or,  if  further  purifi- 
cation is  desirable,  it  may  be  dissolved  in  water,  acidified  with 
hydrochloric  acid,  and  the  filtered  and  dried  product  dissolved 
in  ether  and  shaken  with  animal  charcoal,  when,  upon  evapo- 
ration, the  pure  isonitrosoguaiacol  crystallizes.  Of  the  above 
methods  of  preparation  the  latter  is  much  to  be  preferred, 
although  Rupe  mentions  it  unfavorably  because  of  poor  yields 
and  impure  products.  He  advocates  the  use  of  ethyl  nitrite 
in  a  closed  tube ;  but  on  trial  we  were  unable  to  obtain  the 
quantitative  yields  which  he  reports  and  the  seventy  per  cent 
yield  which  the  amyl  nitrite  gives,  makes  this  method  quite 
equal,  in  efficiency  as  well  as  purity  of  reaction-product,  to 
the  other  more  tedious  process. 

The  silver  salt,  formed  from  the  sodium  salt  by  treating  the 
aqueous  solution  with  a  slight  excess  of  silver  nitrate,  comes 

*  Monatsh.  Chem.,  xix,  536.  t  Ber.  Dtsch.  chem.  Ges.,  xvii,  399. 


THE  ETHERS  OF  ISONITROSOGUAIACOL.          307 

down  as  a  brown  gelatinous  precipitate,  which  becomes  crystal- 
line on  gently  warming,  or  may  be  obtained  in  crystalline  form 
at  once  by  heating  the  separate  solutions  to  50°  C.  before  mix- 
ing. It  is  a  very  unstable  salt,  the  dry  product  decomposing 
with  a  very  gentle  heating. 

Isonitrosoguaiacol  Benzoyl  Ether. 

The  sodium  salt  formed  as  above  was  dissolved  in  as  little 
water  as  possible  and  four  or  five  tunes  its  volume  of  alcohol 
added.  This  solution  was  thoroughly  shaken  with  a  slight 
excess  of  benzoyl  chloride,  added  drop  by  drop.  The  reaction 
is  immediate  and  the  ether  soon  begins  to  come  out  of  the 
solution  in  almost  pure  condition.  Recrystallized  from  alco- 
hol, it  separates  in  straw-colored,  branching  crystals,  which 
melt  sharply  at  188°  C.  when  "  dipped "  for  ten  seconds. 
Heated  gradually  from  normal  temperatures,  it  begins  to  de- 
compose at  175°  C.  and  liquefies  at  185°-188°  C.,  the  temper- 
ature depending  on  the  rapidity  with  which  heat  is  applied. 

Fractional  crystallization  from  alcohol  or  other  solvents  did 
not  essentially  change  the  melting-point,  nor  were  any  differ- 
ent phenomena  observed  when  the  isonitrosoguaiacol  was  made 
by  the  acid  reaction.  This  ether  dissolved  readily  in  chloro- 
form and  glacial  acetic  acid,  much  less  in  benzene  and  ligroin, 
and  is  practically  insoluble  in  ether  and  carbon  disulphide. 
On  analysis : 

0.1100  gram  of  the  substance,  dried  over  H2S04,  gave  0.2632 
gram  C02,  and  0.0428  gram  H20. 

0.1205  gram  of  the  substance  gave  5.45  cm8  K  at  15°  C.  and 
772  mm.  pressure. 

Calculated  for  ™       d 

C14HU04N.  Foun<L 

C  65.34  65.26 

H  4.31  4.32 

N  5.46  5.38 

Isonitrosoguaiacol  Benzoyl  Ether  Dibromide. 

Three  grams  of  the  benzoyl  ether  were  dissolved  in  25  cm3 
of  chloroform  and  2  grams  of  bromine  were  added,  the  mix- 


308          THE  ETHERS  OF  ISONITROSOGUAIACOL. 

ture  being  kept  cool  by  running  water.  The  solution  was 
allowed  to  stand  for  from  one  to  two  hours  and  then  to  evapo- 
rate spontaneously.  The  dibromide  is  left  behind  as  a  light- 
brown  substance.  Attempts  to  crystallize  from  glacial  acetic 
acid  proved  unsatisfactory  since  the  boiling  solvent  caused  a 
decomposition  of  the  ether;  nor  did  a  solution  in  fuming 
nitric  acid,  precipitated  by  water,  give  a  pure  white  product. 
The  best  results  were  obtained  by  washing  with  warm  dilute 
alcohol  when  the  product  becomes  yellowish-white,  and  melts 
at  153°-154°  C.,  browning  considerably  above  140°  C.  An 
analysis  of  the  substance  purified  in  this  manner  shows  a  low 
percentage  of  bromine. 

0.1052  gram  of  the  substance,  dried  over  H2S04,  gave  0.1547 
gram  C02,  and  0.0284  gram  H20. 

0.0509  gram  of  the  substance  gave  0.0449  gram  AgBr. 

Calculated  for  Found. 

C  *  40.29  40.10 

H  2.66  3.00 

Br  38.34  37.52 

Monobromisonitrosoguaiacol  Benzoyl  Ether. 

By  boiling  the  white  dibromide  with  60  per  cent  alcohol  for 
three-quarters  of  an  hour,  the  halogen  acid  and  a  bright-yel- 
low monobromine  compound  are  formed.  This  ether  crystal- 
lizes from  boiling  alcohol  in  the  characteristic  prismatic  form 
of  such  bodies  and  melts  with  decomposition  at  178°  C. 
Repeated  recrystallizations  did  not  change  this  melting-point, 
and  no  indication  was  noted  suggesting  the  possibility  of  a 
mixture.  On  analysis: 

0.1317  gram  of  the  substance,  dried  over  H2S04,  gave  0.2410 
gram  C02,  and  0.0368  gram  H20. 

0.2101  gram  of  the  substance  gave  0.1172  gram  AgBr. 

Calculated  for  ,,       . 

CMH1004NBr.  Found' 

C  49.99  49.91 

H  3.00  3.11 

Br  23.79  23.75 


XXXIX 

THE  CONSTITUTION  OF  THE  AMMONIUM 
MAGNESIUM  ARSENIATE   OF  ANALYSIS. 

BY  MARTHA  AUSTIN.* 

THE  striking  analogy  between  the  phosphates  and  the  arse- 
niates  led  Levolf  to  undertake  the  separation  of  an  ammonium 
arseniate  corresponding  to  the  ammonium  magnesium  phos- 
phate, the  composition  of  which  Berzelius  had  given.  Levol 
states  that  ammonium  magnesium  arseniate  of  the  composition 
NH4MgAsO4 .  10H2O  is  obtained  by  adding  a  solution  of  a 
double  ammonium  magnesium  salt  to  arsenic  acid,  and  that  it 
is  a  salt  possessing  about  the  same  degree  of  solubili ty  in  water, 
in  ammoniacal  water,  and  hi  ammoniacal  water  containing 
magnesium  salt,  as  the  corresponding  phosphate.  Further,  he 
found  that  by  heating  this  salt  to  red  heat  after  carefully 
drying,  magnesium  pyroarseniate  was  given,  from  which 
arsenic  can  be  estimated  readily. 

WachJ  and  H.  Rose  §  obtained  the  ammonium  magnesium 
arseniate  containing  six  molecules  of  water  of  crystallization  by 
precipitating  arsenic  acid  with  magnesia  mixture  and  then  add- 
ing an  excess  of  ammonia,  and  by  drying  at  100°  C.  were  able 
to  estimate  the  arsenic  present  as  the  ammonium-magnesium 
arseniate  containing  one-half  molecule  of  water.  This  method 
seemed  to  offer  an  advantage  over  the  method  of  estimation  as 
the  pyroarseniate,  for  results  obtained  below  the  theoretical 
amount  of  arsenic  present  gave  rise  to  the  suspicion  that 
during  ignition  arsenic  was  reduced  by  the  ammonia  driven  off. 
Rose  attempted  hi  another  way  to  avoid  this  loss  by  igniting  in 
a  current  of  oxygen ;  and  later  Reichelt||  ignited  the  residue 

*  From  Am.  Jour.  Sci.,  ix,  65.  t  Ann.  Chim.  in,  xvii,  501  (1846). 

t  Schweigger,  Jour.  f.  Ch.  u.  Pbys.,  lix,  297. 

§  Ann.  Phys.,  clii,  20  (1849).  ||  Zeitschr.  anal.Chem.,  xx,  89. 


310  CONSTITUTION  OF  THE  AMMONIUM 

after  carefully  saturating  it  with  ammonium  nitrate  and  nitric 
acid,  and  drying  at  100°  C.  Rammelsberg  *  believed  that 
it  was  safer  to  ignite  after  drying  at  120°  C.  because 
drying  at  100° -110°  C.  caused  a  loss  of  ammonia  before  ignition. 
Kaiser  f  dried  the  residue  in  a  current  of  air. 

A  second  source  of  error  discussed  by  H.  Rose4  Fresenius, 
and  others  is  due  to  the  solubility  of  the  ammonium  magnesium 
arseniate  in  water,  in  ammoniacal  water,  and  ammoniacal  water 
containing  magnesium  salts.  Wood  §  attempted  to  avoid 
this  by  precipitating  the  ammonio-magnesium  arseniate  with 
an  alcoholic  magnesia  mixture,  and  by  washing  the  precipitate 
with  an  alcoholic  solution,  by  weighing  the  residue  ignited  in 
a  crucible  held  in  a  second  protecting  crucible,  after  treating 
with  nitric  acid  and  ammonium  nitrate,  he  obtained  concordant 
results.  Brauner  ||  followed  this  method  with  success. 

As  has  been  shown  by  Neubauer,l[  ammonium  chloride 
tends  to  form  a  phosphate  of  magnesium  too  rich  in  ammonia, 
and  in  papers  **  from  this  laboratory  the  tendency  of  ammonia 
and  ammonium  chloride  to  influence  the  proportion  of  the 
metal  in  the  ammonium  phosphate  of  manganese,  of  magnesium 
and  of  other  metals  in  the  second  group  of  Mendele*eff  has 
been  pointed  out.  The  marked  similarity  as  to  behavior 
between  the  phosphates  and  the  arseniates  led  to  the  investi- 
gation of  the  constitution  of  the  ammonium  magnesium 
arseniate  under  the  usual  conditions  imposed  in  analysis,  and, 
further,  to  the  effect  of  ammonium  chloride  on  the  salt  of 
ideal  constitution. 

In  the  table  which  follows  are  recorded  a  set  of  qualitative 
tests,  hi  which  the  hot  filtrates  were  tested  by  hydrogen 
sulphide  in  presence  of  hydrochloric  acid.  These  tests  were 
made  to  show  under  what  conditions  of  volume,  given  amounts 
of  arsenic  acid  can  be  entirely  removed  by  magnesia  mixture 

*  Ber.  Dtsch.  chem.  Ges.,  vii,  544.  t  Zeitschr.  anal.  Chem.,  xiv/250. 

t  Zeitschr.  anal.  Chem.,  iii,  206.  §  Am.  Jour.  Sci.  Ill,  vi,  368. 

||  Zeitschr.  anal.  Chem.,  xvi,  57. 

1  Zeitschr.  anorg.  Chem.,  ii,  45 ;  iv,  251 ;  x,  60.  Zeitschr.  angew.  Chem., 
1896,  435.  Jour.  Am.  Chem.  Soc.,  xiv,  289. 

**  Am.  Jour.  Sci.,  vi,  233 ;  vii,  187 ;  viii,  206.    This  volume,  pp.  121, 190, 252. 


MAGNESIUM  ARSENIATE  OF  ANALYSIS. 


311 


from  ammoniacal  solutions  alone,  or  from  ammoniacal  solutions 
containing  ammonium  chloride.  The  ammonium  chloride  for 
this  work  was  carefully  purified  by  heating  it  to  boiling 
temperature  in  concentrated  solution  —  1  grm.  to  3  cm3  — 
with  ammonia  in  slight  excess.  The  magnesia  mixture  was 
prepared  by  dissolving  one  hundred  and  ten  grams  of  the 
crystallized  magnesium  chloride  in  a  small  volume  of  water, 
filtering  and  adding  to  it  fifty-eight  grams  of  ammonium 
chloride  in  solution,  purified  by  adding  bromine  water  and 
bleaching  with  ammonia,  filtering,  and  then  diluting  to  a 
volume  of  two  liters,  adding  enough  ammonia  — 10  cm3  —  to 
make  this  solution  smell  distinctly  of  ammonia. 

The  results  (1)  to  (3)  recorded  in  the  table  were  obtained 
by  precipitating  the  arsenic  present  in  solution  by  means  of 
the  magnesia  mixture  prepared  as  described  and  rendering  the 

TABLE  I. 


Exp. 

Volume. 

Arsenic 
present 
in  terms 
ofA8a05. 

Magnesia 
mixture. 

NH4OH. 

NH^Cl. 

Indications  of 
arsenic  in 
the  filtrate 
byH2S. 

cm* 

grm. 

cm3 

cm' 

grm. 

(1) 

100 

0.05 

20 

Slight  excess. 

t 

None. 

(2) 

200 

0.05 

20 

Slight  excess. 

m 

None. 

(3) 

300 

0.05 

120 

Slight  excess. 

.  . 

i  Present. 

H20 

ii  Present. 

i"20 

m  None. 

8 

200 
300 

0.5 
0.5 

50 
50 

Slight  excess. 
Slight  excess. 

None. 
Present. 

(6) 

300 

0.5 

50 

2 

Present. 

(7) 

200 

0.5 

60 

4 

Present. 

(8) 

200 

0.5 

30 

2 

Present. 

(9) 

200 

0.5 

130 

4 

1  Present. 

ii  10 

ii  None. 

(10) 

100 

0.5 

130 

2 

1  Present. 

ii  10 

ii  None. 

(11) 

130 

0.5 

60 

Excess. 

10 

Present. 

(12) 

200 

0.5 

50 

Excess. 

10 

Present 

(13) 

150 

0.5 

75 

Excess. 

10 

None. 

(14) 

250 

0.5 

76 

Excess. 

10 

Present. 

(15) 

300 

0.5 

75 

2 

10 

Trace. 

(16) 

400 

0.5 

76 

2 

10 

Present. 

(17) 

285 

0.5 

100 

Excess. 

10 

None. 

(18) 

215 

0.5 

100 

Excess. 

20 

None. 

(19) 

335 

0.6 

100 

Excess. 

60 

Present. 

(20) 

360 

0.5 

125 

Excess. 

60 

Present. 

(21) 

360 

0.5 

150 

Excess. 

60 

None. 

312 


CONSTITUTION  OF  THE  AMMONIUM 


solution  distinctly  ammoniacal.  After  standing  until  the 
precipitate  subsided,  the  solution  was  filtered  and  tested  for 
arsenic  by  hydrogen  sulphide.  Within  certain  limits  of 
volume  all  the  arsenic  is  removed  from  solution  by  the  magnesia 
mixture.  In  case  of  (3)  two  additional  amounts  of  the 
precipitant  had  to  be  made  in  order  to  remove  the  arsenic,  the 
solution  being  ammoniacal.  Results  (4)  to  (10)  of  the  table 
were  obtained  in  the  same  manner  as  (1)  to  (3),  the  amount 
of  arsenic  acid  present  being  increased  ten  times.  In  (9)  and 
(10),  where  the  smaller  amounts  of  magnesia  mixture  were 
used,  additional  portions  had  to  be  added  before  all  the  arsenic 
was  removed  from  solution.  When  the  larger  amounts  of 
ammonia  were  used  the  results  did  not  seem  to  be  influenced, 
(7)  and  (9).  The  results  recorded  in  (11)  to  (21)  were 
obtained  by  precipitating  the  arsenic  in  solution  hi  presence  of 
ammonium  chloride  by  the  magnesia  mixture.  It  is  evident 
that  the  presence  of  ammonium  chloride  causes  some  arsenic 
to  be  dissolved  and  further  that  this  solvent  effect  is  overcome 
by  the  magnesia  mixture  added  in  larger  amounts,  even  when 
the  ammonium  chloride  present  amounts  to  sixty  grams  in 
weight,  as  in  (21). 

In  order  to  find  how  much  arsenic  is  dissolved  from  the 
ammonium  magnesium  arseniate  once  precipitated  so  that  no 
arsenic  is  left  in  solution,  experiments  were  made  in  which  the 
arsenic  in  the  filtrate  was  weighed  after  precipitating  by  hydro- 
gen sulphide  in  hot  acid  solution,  filtering  off  on  asbestos 
under  pressure,  washing  successively  with  water,  alcohol,  car- 
bon disulphide,  alcohol,  and  water,  and  drying  at  100°  C. 

TABLE   II. 


H,O  containing  1  cm3. 
NB4Ofl  in  200  cm». 

NH4OH 

(sp.  gr.  0.96). 

As805 
digested. 

As-jOrj  found 
as  As2S8. 

grm. 

cm8 

grm. 

grm. 

100 

0.5 

0.0019 

100 

0.5 

0.0026 

100 

•    •    . 

0.5 

0.0003 

100 

•   •    • 

0.5 

0.0005 

10 

. 

0.5 

0.0002 

10 

.  .  . 

0.5 

0.0004 

MAGNESIUM  ARSENIATE  OF  ANALYSIS.          313 

It  is  evident  from  these  qualitative  tests  that,  so  far  as  con- 
cerns the  amount  of  arsenic  dissolved  from  the  ammonium 
magnesium  arseniate,  it  is  safe  to  use  a  faintly  ammoniacal 
wash  water  in  small  amounts  —  less  than  100  cm3  —  to  remove 
traces  of  reagents  from  the  residue  after  it  is  gathered  upon 
the  asbestos  felt.  Usually  25-50  cm3  of  wash  water  were  used 
in  rinsing  off  the  precipitate  hi  the  experiments  about  to  be 
given  below. 

The  solution  of  arsenic  employed  in  this  work  was  prepared 
by  dissolving  ten  grams  of  pure  arsenious  oxide,  carefully 
resublimed,  in  a  platinum  dish  in  an  excess  of  pure  nitric  acid 
and  evaporating  on  the  waterbath  to  dryness,  dissolving  the 
arsenic  acid  produced  in  water  and  diluting  to  a  volume  of 
one  liter  in  a  standard  flask. 

Definite  portions  of  this  solution  were  drawn  from  a  burette 
into  a  platinum  dish  and  precipitated  with  magnesia  mixture, 
prepared  as  described,  in  the  proportion  shown  necessary  by 
the  qualitative  tests  for  the  complete  removal  of  the  arsenic 
from  solution,  and  the  solution  was  made  distinctly  ammoni- 
acal. The  precipitate,  dissolved  in  hydrochloric  acid  in  slight 
excess,  was  brought  down  again  by  ammonia  in  distinct  excess. 
After  standing  until  the  precipitate  had  completely  subsided, 
the  precipitate  was  gathered  on  an  asbestos  felt  in  a  perforated 
platinum  crucible,  making  use  of  the  filtrate  to  remove  the 
last  portions  of  the  precipitate  to  the  felt,  before  washing  any 
reagents  from  it  on  the  felt  with  faintly  ammoniacal  water. 
After  carefully  drying,  the  residue  was  ignited.  The  results 
shown  hi  section  A  of  Table  III  fall  so  far  below  the  theory  for 
magnesium  pyroarseniate  that  it  seems  evident  that  the  arsen- 
iate shows  a  tendency  here  to  form  a  salt  richer  hi  ammonia 
than  the  ideal  MgNH4AsO4,  and  yielding  on  ignition  some 
meta-arseniate  instead  of  the  normal  pyroarseniate.  In  section 
B  the  magnesium  ammonium  arseniate  was  precipitated  by  add- 
ing to  the  solution  of  arsenic  acid  magnesia  mixture  contain- 
ing no  free  ammonia  in  the  proportion  necessary  (50  cm3)  to 
remove  the  arsenic  from  solution,  and  then  making  the  solution 
distinctly  ammoniacal.  After  the  precipitate  had  subsided,  it 


314 


CONSTITUTION  OF  THE  AMMONIUM 


was  filtered  off  on  asbestos  under  pressure  in  a  perforated 
platinum  crucible  washed  on  the  felt  with  ammoniacal  water, 
dried  and  ignited.  These  results  are  also  below  the  theory 
for  the  pyroarseniate.  Evidently  the  conditions  here  are  even 
better  for  the  formation  of  the  salt  too  rich  in  ammonia  than 
they  were  in  the  first  case.  No  arsenic  was  found  by  hydro- 
gen sulphide  in  any  case  either  in  the  nitrate  or  in  the  wash 
water  after  acidifying  and  heating. 

TABLE  IIL 


A. 

Mg2As207  corresponding  to  A&jO-. 

As205  found 
byBLjSin 
the  filtrate. 

Taken. 

Found. 

Error. 

grm. 
0.7843 
0.7843 

grm. 
0.7800 
0.7794 

grm. 
0.0043- 
0.0049- 

None. 
None. 

B. 

0.7843 
0.7843 

0.7772 
0.7769 

0.0071- 
0.0074- 

None. 
None. 

In  section  A  of  Table  IV  the  results  recorded  were  obtained 
by  precipitating  definite  portions  of  arsenic  acid  drawn  from 
a  burette  into  a  platinum  dish  with  the  distinctly  ammoniacal 
magnesia  mixture  in  proper  proportion,  and  afterward  adding 
a  little  more  ammonia,  filtering  off  on  asbestos  under  pressure 
in  a  perforated  platinum  crucible  as  soon  as  the  precipitate 
subsided,  washing  off  on  the  felt  with  ammoniacal  water  any 
of  the  reagents  left  on  the  precipitate  in  transferring  it  to  the 
felt  by  using  the  filtrate,  drying,  and  igniting.  No  arsenic 
was  found  in  any  case  in  the  filtrate  or  in  the  wash  water  by 
hydrogen  sulphide.  The  conditions  of  precipitation  here 
prove  to  be  such,  as  the  results  show,  that  the  salt  of  ideal 
constitution  is  formed.  Comparing  these  results  with  those 
of  Table  III,  it  seems  that  the  conditions  under  which  the  salt 
of  ideal  composition  is  formed  are  such  that  at  the  moment 


MAGNESIUM  ARSENIATE  OF  ANALYSIS. 


315 


and  in  the  locality  of  precipitation  the  amount  of  magnesium 
chloride  in  a  certain  volume  of  solution  must  be  large  in  pro- 
portion to  the  amount  of  ammonia  present  as  the  chloride 
and  the  hydroxide ;  otherwise  the  ammonium  of  the  ammonium 
arseniate,  naturally  formed  first,  does  not  suffer  sufficient  dis- 
placement by  magnesium  to  produce  the  normal  ammonium 
magnesium  arseniate.  At  all  events,  the  treatment  applied  in 
the  experiments  of  A  resulted  practically  in  the  complete  pre- 
cipitation and  in  the  production  of  a  precipitate  of  nearly 
ideal  constitution. 

TABLE  IV. 


A. 

Exp. 

Mg2As2O7  corresponding  to  AsjO6. 

Magnesia 
mixture. 

NH4C1. 

Taken. 

Found. 

Error. 

R 

| 

grm. 

0.7843 
0.7843 
0.7843 
0.7843 

grm. 

0.7830 
0.7849 
0.7841 
0.7843 

gnu. 
0.0013- 
0.0006+ 
0.0002- 
0.0000 

cm8 
50 
60 
50 
50 

grm. 

B. 

(5) 
(6) 

11! 

(9) 
(10) 

(11) 
(12) 

0.7843 
0.7843 
0.7843 
0.7843 
0.7843 
0.7843 
0.7843 
0.7843 

0.7763 
0.7762 
0.7832 
0.7838 
0.7784 
0.7810 
0.7849 
0.7846 

0.0080- 
0.0081- 
0.0011- 
0.0005- 
0.0059- 
0.0033- 
0.0006+ 
0.0003+ 

75 
75 
100 
100 
100 
100 
150 
150 

10 
10 
10 
10 
20 
20 
60 
60 

In  section  B  of  the  same  table  the  results  recorded  show  the 
effect  of  increased  amounts  of  ammonium  chloride  on  the 
constitution  of  the  ammonium  magnesium  arseniate.  Enough 
magnesia  mixture  was  used  in  each  case  to  remove  the  arsenic 
completely  from  solution.  The  precipitates  were  gathered  on 
asbestos  in  a  perforated  platinum  crucible  and  treated  like 
those  described  in  section  A  of  the  table.  It  is  evident  as 
shown  by  (5)  and  (6)  of  the  table  that  ammonium  chloride 
causes  a  replacement  of  some  of  the  metal  by  ammonia  in  the 


316  AMMONIUM  MAGNESIUM  ARSENIATE. 

ammonium  magnesium  arseniate  (to  form,  possibly,  a  salt  of 
the  constitution  Mg(NH4)4AsO4)  though  the  solvent  effect 
of  the  ammonium  chloride  is  overcome  by  the  addition  of  a 
sufficient  amount  of  magnesia  mixture.  No  arsenic  appeared 
in  the  filtrates,  and,  further,  experiments  (7)  and  (8)  show 
that  increasing  the  amount  of  magnesia  mixture  present  will 
cause  the  formation  of  the  salt  of  ideal  constitution  even  in 
presence  of  considerable  amounts  of  ammonium  chloride. 
Indeed  this  is  possible  where  as  large  an  amount  as  sixty 
grams  of  the  salt  is  present,  as  results  (11)  and  (12)  show. 
Obviously,  ammonium  chloride  in  any  amount  above  what  is 
required  for  the  magnesia  mixture  tends  to  dissolve  the 
precipitate,  but  this  solvent  effect  may  be  neutralized  by 
increasing  the  amounts  of  magnesia  mixture  even  though  the 
precipitate  formed  is  richer  in  ammonia  than  the  ideal  salt. 

In  no  one  of  the  many  precipitates  tested  by  silver  nitrate 
for  included  chlorides  was  more  than  an  inappreciable  trace 
found. 

Evidently,  when  ammoniacal  magnesia  mixture,  amounting 
to  about  thirty  cubic  centimeters  in  excess  of  the  theoretical 
amount  necessary  to  precipitate  all  the  arsenic  as  the  ammo- 
nium magnesium  arseniate,  is  added  to  the  faintly  acid 
solution  of  arsenic  acid  (carrying  no  ammonium  salts)  in  a 
volume  not  exceeding  two  hundred  cubic  centimeters,  the 
precipitate  appears  to  fall  in  ideal  condition.  If  the  precipi- 
tated salt  is  transferred  to  the  filtering  crucible  by  the  aid  of 
portions  of  the  filtrate  used  as  the  washing  liquid  and  finally 
washed  on  the  asbestos  with  about  twenty-five  cubic  centimeters 
of  faintly  ammoniacal  water  —  an  amount  which  is  quite 
sufficient  after  the  transfer  has  been  made  —  no  arsenic  gets 
into  solution.  The  weight  of  the  carefully  dried  and  ignited 
pyroarseniate  indicates  with  accuracy  the  amount  of  arsenic 
present. 


XL 

ON  THE  ESTIMATION  OF  THALLIUM  AS  THE 
ACID  AND  NEUTRAL  SULPHATES. 

BY  PHILIP  E.  BROWNING.* 

CEOOKES  t  has  shown  that  the  salt  obtained  by  heating 
thallous  chloride  with  sulphuric  acid  until  the  excess  of  the 
latter  is  expelled  and  then  raising  the  heat  to  redness  has  the 
constitution  of  a  neutral  sulphate, 

He  also  found  that  continued  heating  did  not  result  in 
any  essential  loss  of  weight,  and  suggested  the  possibility 
of  applying  this  method  of  treatment  to  the  estimation  of 
thallium. 

CastanjenJ  in  a  recent  paper  discusses  thoroughly  the 
compounds  of  thallium  and  confirms  essentially  the  statements 
of  Crookes  in  regard  to  the  neutral  sulphate,  adding,  however, 
the  observation  that  on  strong  ignition  in  the  air  this  salt 
tends  to  lose  sulphuric  acid.  He  also  mentions  in  the  same 
paper  the  acid  sulphate,  and  states  that  on  heating  it  first 
melts  and  on  continued  heating  gives  off  sulphuric  acid, 
leaving  the  neutral  sulphate. 

The  work  to  be  described  in  this  paper  was  undertaken  to 
determine  under  what  conditions  the  formation  of  these  salts 
may  be  applied  to  the  estimation  of  thallium.  For  the  work  a 
solution  was  made  by  dissolving  a  given  amount  of  the  nitrate 
in  water  and  making  up  to  a  liter.  The  value  of  the  solution 
was  determined  by  precipitating  measured  and  weighed 
amounts  of  this  solution  both  as  the  iodide  and  chromate,  as 
described  in  a  previous  paper.  §  Closely  agreeing  results  by 

*  From  Am.  Jour.  Sci.,  ix,  137.  t  Chem.  News,  viii,  243. 

t  Jour,  prakt.  Chem.,  cii,  131. 

§  Am.  Jour.  Sci.,  yiii,  460.    This  volume,  p.  300. 


318 


ESTIMATION  OF  THALLIUM  AS  THE 


both  methods  were  taken  as  the  standard.  Measured  amounts 
of  this  solution  were  drawn  from  a  burette  into  weighed 
platinum  crucibles,  and  the  weight  taken  as  a  check  on  the 
burette  reading.  To  the  solution  in  the  crucible  a  few  drops 
of  sulphuric  acid  were  added  and  the  water  removed  by 
evaporation  over  a  steam  bath.  The  crucible  was  then  removed 
to  a  radiator,  consisting  of  a  conical  iron  cup,  and  heated 
at  a  temperature  ranging  from  220°  C.  to  240°  C.,  until 
fuming  ceased  and  the  weight  after  half-hour  periods  of 
heating  remained  constant.  The  crucibles  were  placed  in  the 
radiator  upon  a  pipe  stem  triangle  so  that  they  were  about  5 
cm.  from  the  bottom,  which  was  heated  at  low  redness.  A 
thermometer,  hung  so  that  the  bulb  occupied  the  same  position 
as  the  crucible,  gave  the  reading  mentioned  above. 

As  will  be  seen,  the  results  obtained  by  this  treatment  agree 
closely  with  the  calculated  amounts  of  acid  sulphate  which 
should  be  formed.  In  several  experiments  this  salt  was  dis- 
solved in  water  and  the  sulphuric  acid  present  in  combination 
precipitated  by  barium  nitrate.  The  results  obtained  agreed 
closely  with  the  formula  of  the  acid  sulphate  of  thallium. 
Having  obtained  by  the  method  described  the  acid  sulphate, 
the  crucibles  were  removed  and  heated  carefully  over  a  free 
flame  to  low  redness,  when,  after  a  considerable  Devolution  of 
sulphuric  acid  fumes,  the  weight  again  became  constant,  and 
the  results  showed  a  condition  closely  approximating  to  that  of 
the  neutral  sulphate.  In  several  of  these  experiments  the 


Exp. 

T1HS04 
calculated. 

T1HS04 
found. 

Error. 

T12S04 
calculated. 

T12S04 
found. 

Error. 

grin. 

grm. 

grin. 

grm. 

grm. 

grm. 

(1) 

0.1605 

0.1596 

0.0009- 

0.1344 

0.1346 

0.00024- 

(2) 

0.1611 

0.1608 

0.0003- 

0.1349 

0.1346 

0.0003- 

(3) 

0.1608 

0.1608 

0.0000 

0.1347 

0.1352 

0.0005-f 

(4) 

0.1  6T2 

0.1600 

0.0012- 

0.1350 

0.1346 

0.0004- 

5) 

0.1602 

0.1596 

0.0006- 

0.1341 

0.1346 

0.0005-j- 

(6) 

0.1608 

0.1596 

0.0012- 

. 

. 

. 

(7) 

0.1617 

0.1604 

0.0013- 

t  t 

4 

(8) 

0.1608 

0.1592 

0.0016- 

0.1347 

0.1358 

0.00114- 

(9) 

0.1609 

0.1590 

0.0019- 

0.1348 

0.1346 

0.0002- 

ACID  AND  NEUTRAL  SULPHATES.  319 

sulphuric  acid  present  in  combination  was  determined,  and 
showed  amounts  closely  agreeing  with  the  constitution  of  the 
neutral  sulphate. 

These  results  would  seem  to  show  that  thallium  may  be 
estimated  either  as  the  acid  sulphate  or  as  the  neutral  sulphate 
by  careful  attention  to  the  proper  conditions  of  temperature. 


XLI 

THE    SEPARATION     AND    DETERMINATION    OF 
MERCURY  AS  MERCUROUS   OXALATE. 

BY  C.  A.  PETERS.* 

IT  is  stated  in  the  literature!  that  oxalic  acid,  neutral  and 
acid  oxalates  of  the  alkalies,  precipitate  mercurous  salts,  and 
that  oxalic  acid  and  the  double  oxalates  of  potassium  produce 
no  precipitate  with  mercuric  chloride  solution.  Starting  with 
these  facts,  the  attempt  was  made  to  estimate  mercurous  salts : 
volumetrically,  by  precipitating  with  ammonium  oxalate  and 
determining  the  oxalic  acid  by  potassium  permanganate ;  and 
gravimetrically  by  direct  weighing  of  the  precipitate. 

The  Volumetric  Estimation. 

The  mercurous  nitrate  solution  used  was  standardized  by 
the  battery,  and  contained  about  12  grm.  of  metallic  mercury 
to  the  liter.  To  obviate  the  tendency  of  the  mercury  salt 
to  break  down  and  form  basic  salts  J  upon  the  addition  of  a 
large  amount  of  water  if  no  nitric  acid  is  added,  the  solution 
was  prepared  in  the  following  manner.  About  20  grm.  of 
mercurous  nitrate  were  ground  in  a  mortar,  transferred  to  a 
flask,  and  200-300  cm3  water  added.  After  shaking  well, 
the  solution  was  filtered  and  the  nitrate  diluted  to  one  liter. 
Five  cubic  centimeters  of  this  solution  when  precipitated  with 
a  sodium  chloride  solution  gave  a  nitrate  from  which  only  a 
very  slight  darkening  in  color  could  be  obtained,  even  upon 

*  From  Am.  Jour.  Sci.,  ix,  401. 

t  Rose-Finkener,  Handbuch  der  analytischen  Chemie,  i,  319. 

J  Graham-Otto,  Handbuch,  iii,  1102. 


SEPARATION  OF  MERCURY. 


321 


several  hours'  standing,  when  treated  with  hydrogen  sulphide, 
thus  showing  the  absence  of  a  mercuric  salt. 

A  solution  made  in  the  above  manner  had  not  changed 
its  standard  after  a  period  of  eight  weeks.  The  potassium 
permanganate  solution  (approximately  ^)  was  standardized 
against  lead  oxalate. 

It  was  first  attempted  to  estimate  the  mercurous  salts  as 
follows.  The  mercurous  oxalate  was  precipitated  cold  by 
means  of  ammonium  oxalate,  stirred  well,  and  allowed  to 
settle,  the  completion  of  the  precipitation  being  determined 
by  addition  of  more  ammonium  oxalate.  The  precipitate  was 
collected  on  asbestos,  washed  once  or  twice  with  cold  water, 
and  (still  in  the  crucible)  treated  in  a  beaker  with  5  cm8  of 
strong  hydrochloric  acid.  To  the  solution  diluted  to  100-200 
cm3  1  grm.  of  a  manganous  salt  was  added,  and  the  oxalic 
acid  was  titrated  with  permanganate  at  the  ordinary  tem- 
perature of  the  room.  The  end  color  was  not  stable  and 
was  hard  to  determine.  Three  experiments,  using  0.1217  of 
mercury  hi  form  of  mercurous  nitrate,  gave  plus  errors  of 
0.0011  grm.,  0.0017  grm.  and  0.0028  grm.  respectively,  or  1.5 
per  cent.  The  precipitate  when  dissolved  in  sulphuric  acid 


Excess  of 

ammonium 

Exp. 

Hg  taken  as 

Hg2(N08), 

Hg  found. 

oxalate 
approxi- 

MnCl, .  4H2O. 

H2S04 

Error 
asHg. 

mately^. 

' 

A. 

grm. 

grm. 

cm8 

cm» 

grm. 

cm3 

grm. 

(1) 

0.1825 

0.1823 

0.90 

5 

0.5 

0.0002- 

(2) 

0.1217 

0.1218 

0.93 

5 

0.6 

0.0001+ 

(3) 

0.1217 

0.1206 

0.99 

5 

0.5 

0.0011- 

(4) 

0.1217 

0.1210 

4.05 

5 

0.5 

0.0007- 

(5) 

0.3042 

0.3034 

4.97 

5 

0.6 

.  . 

0.0008- 

B. 

(6) 

0.1217 

0.1220 

0.93 

t 

5 

0.0003+ 

(7) 

0.1217 

0.1211 

0.97 

, 

g 

5 

0.0006- 

(8) 

0.1825 

0.1827 

0.89 

5 

0.0002+ 

(9) 
(10) 

0.3042 
0.1217 

0.3040 
0.1202 

0.82 
4.10 

* 

• 

5 
5 

0.0002- 
0.0015- 

TOL.  II.  —  21 


322          SEPARATION  AND  DETERMINATION  OF 

and  titrated  gave  no  better  results.  To  obviate  this  difficulty 
the  ammonium  oxalate  solution  was  matched  on  the  per- 
manganate and  the  oxalic  acid  in  the  nitrate  determined. 
The  results  obtained  by  this  method,  given  in  the  preceding 
table,  are  quite  accurate. 

In  the  experiments  recorded  in  section  A  of  the  table  the 
filtrate  was  titrated  in  the  presence  of  hydrochloric  acid  and 
a  manganous  salt,  at  a  temperature  of  20° -40°.*  In  the 
experiments  under  section  B  sulphuric  acid  was  added  and 
the  solution  heated  in  the  usual  manner.  An  excess  of 
ammonium  oxalate,  as  shown  hi  experiments  (4),  (5)  and 
(10),  interferes  in  no  way. 

The  separation  of  the  mercurous  salt  from  small  quantities  of 
mercuric  salts,  by  the  means  of  dilute  nitric  acid  (sp.  gr.  1.15), 
was  next  attempted.  It  is  stated  f  that  mercurous  oxalate  is 
insoluble  in  cold  dilute  nitric  acid,  while  mercuric  oxalate  is 
more  or  less  soluble  in  the  same  reagent.  Before  attempting 
any  separation,  however,  there  are  three  factors  with  reference 
to  the  action  of  the  nitric  acid  which  need  to  be  deter- 
mined—  first,  the  maximum  amount  of  nitric  acid  that  may 
be  present  in  the  titration  of  an  oxalate  without  interference ; 
second,  the  maximum  amount  of  nitric  acid  that  may  be 
present  in  a  precipitation  of  mercurous  oxalate  without  hav- 
ing any  perceptible  solvent  action  upon  the  same  ;  and,  third, 
the  amount  of  mercuric  oxalate  which  will  be  held  in  solution 
by  given  amounts  of  nitric  acid. 

To  determine  the  amount  of  nitric  acid  that  may  be  present 
in  the  titration  of  an  oxalate,  10  cm3  of  ^  ammonium  oxalate 
were  titrated  with  permanganate  at  a  dilution  of  100  cm3 
with  sulphuric  acid  at  80°  C.  The  event  proved  that  10  cm3 
of  nitric  acid  (sp.  gr.  1.15)  may  be  present  without  appearance 
of  interfering  action.  The  maximum  amount  of  nitric  acid 
which  may  be  present  without  action  upon  the  mercurous  salt 
was  determined  as  shown  in  the  following  experiments. 

*  Gooch  and  Peters,  Am.  Jour.  Sci.,  vii,  461, 1899.    This  volume,  p.  222. 
t  Souchay  and  Lenssen,  Ann.  Chem.  (Liebig),  cii,  43. 


MERCURY  AS  MERCUROUS  OXALATE. 


323 


Excess  of 

Hg  taken  as 

Hg2(N03),. 

ammonium 
oxalate 
approximately 

TO' 

HNO, 

"?:$: 

Volume 
at 
precipitation. 

Hg  found. 

Error. 

gnn. 

era.3 

cm8 

cm8 

grin. 

grm. 

0.1122 

1.64 

10 

100 

0.1087 

0.0035- 

0.1122 

2.62 

6 

100 

0.1098 

0.0024- 

0.1122 

1.59 

6 

100 

0.1096 

0.0026- 

f  0.1122 

1.53 

6 

100 

0.1109 

0.0013- 

J  0-1122 

1.40 

5 

100 

0.1134 

0.0012+ 

1  0.1122 

1.50 

5 

100 

0.1115 

0.0007- 

10.1122 

1.70 

4 

100 

0.1116 

0.0006- 

0.1122 

1.59 

8 

200 

0.1096 

0.0026- 

0.1122 

6.62 

8 

200 

0.1111 

0.0011- 

0.1122 

7.72 

8 

200 

0.1108 

0.0014- 

(  0.1122 

2.59 

5 

200 

0.1107 

0.0015- 

\  0.1010 

2.07 

5 

200 

0.1003 

0.0007- 

Working  under  the  conditions  stated  in  the  above  table,  it 
is  plain  that  5  cm3  of  nitric  acid  (sp.  gr.  1.15)  may  be  used 
before  its  solvent  action  is  sufficient  to  interfere  with  the 
accuracy  of  the  process. 

To  determine  the  amount  of  mercuric  salt  that  would  be 
held  up  by  5  cm3  of  nitric  acid  (sp.  gr.  1.15),  the  following 
experiments  were  made. 


Hg  taken  as 
Hg2(N03),. 

HNO3 

Tiff 

Ammonium 
oxalate^ 
in  excess. 

Volume. 

Time  before  precipitation. 

gms. 

cm** 

cm8 

cm8 

0.0335 

4 

0.76 

100 

6  hours. 

0.0095 

4 

1.00 

100 

No  ppt.,  20  hours. 

0.0143 

5 

6.6 

100 

No  ppt.,  20  hours. 

0.0238 

5 

1.5 

100 

Slight  ppt.  20  hours. 

0.0238 

5 

5.5 

100 

26  minutes. 

Five  cm3  of  dilute  nitric  acid  (sp.  gr.  1.15)  will  prevent  the 
precipitation  of  small  amounts  of  mercuric  salt,  10-20  mgrm. 
calculated  as  mercury,  depending  upon  the  amount  of  am- 
monium oxalate  present  in  excess.  This  amount  of  nitric  acid 
has  no  apparent  solvent  action  on  a  precipitate  of  mercurous 
oxalate  under  conditions  already  stated  and  does  not  interfere 


324 


SEPARATION  AND  DETERMINATION  OF 


with  the  titration  of  an  oxalate  by  permanganate  as  already 
shown. 

Carrying  out  the  process  of  separation  of  mercurous  salts 
from  mercuric  salts,  the  precipitation  was  made  as  described  for 
the  estimation  of  mercurous  salts  alone,  excepting  that  nitric 
acid  and  the  mercuric  salt  were  added.  The  experiments  in  A, 
B,  and  C,  of  the  accompanying  table  show  the  amounts  of  mer- 
curic salt  from  which  the  mercurous  oxalate  may  be  separated 
with  2  cm3  of  nitric  acid.  In  experiments  in  section  A  the 
results  are  quite  accurate,  but  an  excess  of  ammonium  oxalate 
tends  to  increase  the  results  a  little  as  seen  in  experiments 
under  B.  An  increase  in  the  amount  of  mercuric  salt  present 
causes  also,  as  shown  in  section  C,  the  results  to  be  a  trifle 
high. 

In  experiments  D,  using  4  cm3  of  nitric  acid  even  in  the 


Hg  taken  as 

JJg^JfOaU. 

Hg(N03), 
present 
calculated 
asHg. 

Ammonium 
oxalate. 
approxi- 
mately^ 
in  excess. 

HNO8 

Volume 
at 
precipi- 
tation. 

Hg  found. 

Error. 

grm. 

gnu. 

cm» 

cm3 

grm. 

grm. 

grm. 

(  0.1217 

0.0067 

0.86 

2 

100 

0.1232 

0.0015+ 

1  0.1217 

0.0067 

0.92 

2 

100 

0.1220 

0.0003+ 

A 

0.1217 

0.0067 

0.97 

2 

100 

0.1218 

0.0001+ 

I  0.1217 

0.0067 

0.90 

2 

100 

0.1224 

0.0007+ 

1  0.1217 

0.0067 

0.91 

2 

100 

0.1213 

0.0004- 

0.1221 

0.0067 

3.92 

2 

100 

0.1237 

0.0016+ 

B 

0.1217 

0.0067 

3.93 

2 

100 

0.1235 

0.0018+ 

0.1242 

0.0067 

8.75 

2 

100 

0.1263 

0.0021+ 

0.1217 

0.0134 

0.87 

2 

100 

0.1230 

0.0013+ 

0.1217 

0.0134 

0.86 

2 

100 

0.1232 

0.0015+ 

;  0.1217 

0.0134 

3.93 

4 

100 

0.1218 

0.0001+ 

" 

0.1217 

0.0134 

3.90 

4 

100 

0.1211 

0.0006- 

T? 

!  0.2244 

0.0067 

1.88 

4 

115 

0.2244 

0.0000 

E  - 

0.2244 

0.0067 

1.91 

4 

130 

0.2240 

0.0004- 

I  0.2244 

0.0141 

2.98 

4 

100 

0.2230 

0.0014- 

0.2244 

0.0141 

2.94 

4 

100 

0.2241 

0.0003- 

i  0.2244 

0.0067 

8.88 

4 

100 

0.2289 

0.0045+ 

0.2244 

0.0067 

8.90 

4 

100 

0.2285 

0.0041+ 

[  0.2424 

0.0067 

1.75 

4 

200 

0.2432 

0.0008+ 

G 

!  0.2424 

0.0067 

2.00 

4 

200 

0.2414 

0.0010- 

i 

'  0.2424 

0.0067 

2.96 

4 

200 

0.2421 

0.0003- 

0.2241 

0.0134 

1.74 

4 

200 

0.2271 

0.0026+ 

u. 

0.2245 

0.0134 

1.81 

4 

200 

0.2256 

0.0011+ 

I 

0.1122 

0.0144 

6.62 

5 

200 

0.1121 

0.0001- 

(0.1122 

0.0240 

6.54 

5 

200 

0.1136 

0.0014+ 

i  0.1122 

0.0240 

6.57 

5 

200 

0.1130 

0.0008+ 

MERCURY  AS  MERCUROUS  OXALATE.  325 

presence  of  an  excess  of  ammonium  oxalate,  the  results  are 
accurate.  In  experiments  E  the  amount  of  mercurous  salt 
was  doubled  and  the  results  are  accurate.  In  EE  the  amount 
of  mercuric  salt  was  also  doubled  and  the  results  are  still 
fairly  accurate ;  but  when  a  large  excess  of  ammonium  oxa- 
late is  present  as  in  experiments  F,  even  with  the  smaller 
amount  of  mercuric  salt,  the  results  are  high.  At  a  dilution 
of  200  cm3  the  results  are  normal  as  seen  in  experiments  G  ; 
but  the  introduction  of  more  mercuric  salt,  as  in  experiments 
H,  causes  a  plus  error. 

Using  5  cm3  of  nitric  acid,  the  larger  amount  of  mercuric 
salt  together  with  a  large  excess  of  ammonium  oxalate,  as 
recorded  in  experiments  K,  the  error  is  raised  a  trifle;  but 
with  a  smaller  amount  of  mercurous  salt,  as  in  experiment  I, 
the  result  is  normal. 

In  precipitating  mercurous  salts  by  ammonium  oxalate  (^) 
it  is  an  easy  matter  to  keep  the  excess  of  the  precipitant 
within  the  limits  of  1  or  two  cm3,  because  the  mercurous  oxa- 
late, when  properly  stirred,  settled  very  rapidly. 

The  Gravimetric  Estimation. 

All  the  conditions  described  above  in  the  volumetric  estima- 
tion of  mercurous  oxalate,  for  the  separation  of  mercurous 
from  mercuric  salts,  may  be  applied  to  the  gravimetric  estima- 
tion of  mercurous  oxalate.  The  precipitate  is  collected  on  a 
weighed  asbestos  filter,  washed  two  or  three  times  with  cold 
water  and  dried  over  sulphuric  acid  to  a  constant  weight. 
Amounts  of  mercurous  oxalate  equivalent  to  0.1217  and  0.2244 
grm.  of  metallic  mercury  dried  to  a  constant  weight  over  sul- 
phuric acid  in  about  15  hours ;  a  larger  amount  equivalent  to 
0.3  grm.  of  metallic  mercury,  required  about  2  days  to  dry  to  a 
constant  weight.  Souchay  and  Lenssen*  state  that  mercurous 
oxalate  breaks  up  at  100° ;  consequently  this  temperature 
cannot  be  used  for  drying.  For  example,  a  precipitate  con- 
taining 0.1122  grm.  mercury  as  the  oxalate  which  when 

*  Ann.  Chem.  (Liebig),  ciii,  308. 


326 


SEPARATION  AND  DETERMINATION  OF 


brought  to  a  constant  weight  at  the  ordinary  temperature  over 
sulphuric  acid  weighed  0.1371  grm.,  when  heated  weighed  as 
follows  : 

After  7  hours  at  110°,  weight  =  0.1338  grm. 
«     2        "        "  "          0.1328    " 

«     7        ((        «  «          0.1302  « 

The  result  shows  a  loss  of  0.0069  grm.  for  16  hours,  heating, 
and  agrees  with  the  statement  of  Souchay  and  Lenssen. 

The  following  experiments  give  the  results  of  the  gravi- 
metric work,  in  which  the  drying  was  effected,  by  exposure 
for  15  hours  or  less,  at  ordinary  temperatures  over  sulphuric 
acid.  The  larger  amounts  of  nitric  acid  present  in  the  separa- 
tions cause  the  precipitate  to  be  more  granular  and  aid  in  the 
filtering  process. 


Excess  of 

ammonium 

Hg  taken  as 

Hg  present 

oxalate 

HN03 

Volume  at 

Hg2(N03)3 

asHg(N08)2 

present 

precipi- 

Hg found. 

Error. 

approxi- 

1.15). 

tation. 

mately  ^. 

grm. 

grm. 

cm8 

cm8 

cm3 

grm. 

grm. 

0.1217 

. 

2-4 

,  , 

100 

0.1217 

0.0000 

K 

0.1217 

.  .  . 

2-4 

.  . 

100 

0.1217 

0.0000 

0.1122 

.  .  . 

2-4 

.  . 

100 

0.1124 

0.0002+ 

0.1122 

0.0067 

0.93 

2 

100 

0.1130 

0.0008- 

0.1122 

0.0067 

0.93 

2 

100 

0.1112 

0.0010- 

M  "  0.1122 

0.0067 

4.40 

2 

100 

0.1124 

0.0002+ 

N    0.1122 

0.0135 

0.72 

4 

100 

0.1125 

0.0003+ 

0.2244 

0.0071 

1.68 

4 

100 

02253 

0.0009+ 

0.2244 

0.0071 

2.46 

4 

100 

0.2241 

0.0003- 

Tj 

0.2244 

0.0048 

0.54 

4 

200 

0.2248 

0.0004+ 

0.2244 

0.0048 

2.44 

4 

200 

0.2245 

0.0001+ 

In  section  K  are  experiments  showing  the  accuracy  of  the 
process  where  a  mercurous  salt  is  precipitated  in  the  absence 
of  a  mercuric  salt.  A  small  amount  of  mercuric  salt  was 
introduced  in  experiments  L,  and  an  excess  of  ammonium 
oxalate  in  experiment  M,  a  still  larger  amount  of  mercuric 
salt  was  present  in  experiment  N,  and  a  larger  amount  of 
mercurous  salt  in  experiments  under  O.  In  experiments  in 
section  P  a  dilution  of  200  cm3  was  employed  both  with  and 


MERCURY  AS  MERCUROUS  OXALATE.  327 

without  an  excess  of  ammonium  oxalate.  All  the  results  are 
within  reasonable  limits  of  error. 

The  work  may  be  summed  up  briefly  as  follows :  Mercurous 
nitrate  may  be  estimated  volumetrically  by  precipitating  as 
the  oxalate  and  determining  the  excess  of  the  precipitant  with 
permanganate. 

The  precipitated  mercurous  oxalate  may  also  be  estimated 
gravimetrically  by  drying  it  over  sulphuric  acid  and  weighing 
directly. 

In  solutions  containing  2-5  per  cent  dilute  nitric  acid,  sp.  gr. 
1.15,  mercurous  salts  may  be  separated  quantitatively  as  the 
oxalate  from  small  quantities  of  mercuric  salts. 

If  about  0.12  grm.  of  mercury  is  present  as  the  nitrate  in 
100  cm3  of  water,  about  12  per  cent  of  that  amount  of  mercury 
as  the  mercuric  salt  may  be  present  without  interfering  with 
the  accuracy  of  the  estimation,  and  even  20  per  cent  may  be 
present  before  an  appreciable  rise  in  the  result  is  apparent.  If 
the  amount  of  mercurous  salt  present  is  doubled,  the  amount 
of  mercuric  salt  which  may  be  present  is  cut  down  about 
one-half. 


XLII 

THE  TITRATION  OF  MERCURY  BY  SODIUM 
THIOSULPHATE. 

BY  JOHN  T.   NORTON,  JR.* 

ACCORDING  to  J.  J.  Schererf  mercurous  nitrate,  mercuric 
nitrate  and  mercuric  chloride  may  be  estimated  by  direct 
titration  with  sodium  thiosulphate,  Hg2S,  2HgS  .  Hg(NO3)2, 
and  2HgS  .  HgCla  being  the  precipitates  obtained  in  each 
case.  I  have  been  unable  to  obtain  access  to  Scherer's  original 
publication,  but  Suttonf  gives  the  following  very  general 
directions  for  this  process : 

"  (a)  Mercurous  salts.  —  The  solution  containing  the  metal 
as  a  protosalt  only  is  diluted,  gently  heated  and  the  thiosul- 
phate delivered  in  from  the  burette  at  intervals,  meanwhile 
well  shaking  until  the  last  drop  produces  no  brown  color.  The 
sulphide  settles  freely  and  allows  the  end  of  the  reaction 
to  be  easily  seen.  One  cm3  of  the  ^V  normal  solution  of 
thiosulphate  =  0.02  grm.  Hg  or  0.0208  grm.  Hg2O. 

"  (5)  Mercuric  nitrate.  —  The  solution  is  considerably  diluted, 
put  into  a  stoppered  flask,  nitric  acid  added  and  the  thiosul- 
phate cautiously  added  from  the  burette,  vigorously  shaken 
meanwhile,  until  the  last  drop  produces  no  further  precipitate. 
Scherer  recommends  that  when  the  greater  part  of  the  metal 
is  precipitated  the  mixture  should  be  diluted  to  a  definite 
volume,  the  precipitate  allowed  to  settle  and  a  measured 
quantity  of  the  clear  li quid  taken  for  titration ;  the  analysis 
may  then  be  checked  by  a  second  titration  of  the  clear  liquid  if 
needful.  One  cm3  of  ^  normal  thiosulphate  =  0.015  grm.  Hg  or 
0.0162  grm.  HgO. 

*  From  Am.  Jour.  Sci.,  x,  48. 

t  Scherer's  Lehrbuch  der  Chemie,  i,  513. 

J  Volumetric  Analysis,  p.  220. 


THE   TITRATION  OF  MERCURY.  329 

"  (<?)  Mercuric  chloride.  —  With  mercuric  chloride  the  end  of 
the  process  is  not  so  easily  seen.  The  very  dilute  solution  is 
acidified  with  hydrochloric  acid,  heated  nearly  to  boiling,  and 
the  thiosulphate  cautiously  added  so  long  as  a  white  precipitate 
is  seen  to  form ;  any  great  excess  of  the  precipitant  produces 
a  dirty-looking  color.  Filtration  is  necessary  to  distinguish 
the  exact  ending  of  the  reaction.  One  cm3  of  fa  normal 
thiosulphate  =  0.015  grin.  Hg  or  0.0162  grm.  HgO." 

Fresenius  *  gives  practically  the  same  directions,  but  omits 
all  mention  of  that  portion  of  the  process  dealing  with 
mercurous  nitrate. 

In  view,  therefore,  of  the  scant  information  available  on  the 
subject  and  of  the  apparent  difficulty  of  working  the  process 
accurately  according  to  the  directions  given,  an  attempt  was 
made  to  ascertain  whether  the  careful  regulation  of  tempera- 
ture, dilution,  and  amount  of  acid  present  might  not  produce 
beneficial  results. 

That  portion  of  the  process  dealing  with  mercuric  chloride 
was  first  taken  up.  The  mercuric  chloride  used  was  pulverized, 
dried  at  100°  and  its  purity  proved  by  several  determinations 
as  mercuric  sulphide.  The  sodium  thiosulphate  was  made 
up  of  approximately  fa  normal  strength  and  standardized 
on  decinormal  iodine,  which  in  turn  was  titrated  against 
decinormal  arsenious  acid  made  from  pure  resublimed  arsenious 
oxide. 

For  the  action  of  sodium  thiosulphate  upon  the  mercuric 
chloride  Scherer  gives  the  equation, 

3HgCl2  +  2Na2S203  +  3H20  =  2HgS  .  HgCl2  +  2Na2S04  +  4HC1. 

According  to  my  experience,  the  action  results  in  the  formation 
of  a  dense  white  precipitate  which  refuses  to  settle  either  by 
shaking  or  standing,  thus  making  it  impossible  to  fix  the  end 
reaction  by  reading  the  first  drop  of  thiosulphate  which  produces 
no  further  white  precipitate  in  the  solution  containing  the 
mercuric  chloride.  Recourse  must  be  had  therefore  to  filtering. 
By  far  the  quickest  and  neatest  method  is  to  use  the  asbestos 

*  Quantitative  Analysis. 


330 


THE   TITRATION  OF  MERCURY 


filter  deposited  on  a  large  perforated  platinum  cone.*  This 
cone  is  set  in  a  glass  funnel  by  means  of  a  rubber  connector 
and  the  funnel  is  passed  through  the  stopper  of  a  large  side- 
necked  Erlenmeyer  connected  with  an  exhaust  pump.  A  little 
asbestos  fiber  shaken  in  the  liquid  to  be  filtered  was  found  to 
be  very  beneficial  in  preventing  the  precipitate  from  running 
through  the  filter.  In  all  the  following  experiments  the 
thiosulphate  was  run  into  the  solution  containing  the  mercuric 
chloride  in  excess,  the  whole  shaken  up  with  asbestos  fiber, 
filtered  and  the  excess  of  thiosulphate  determined  by  ^ 
normal  iodine.  This  procedure  seems  to  be  far  preferable 
to  attempting  to  catch  the  end  of  the  reaction  by  run- 
ning hi  the  thiosulphate  until  the  last  drop  produces  no 
precipitate.  In  the  experiments  shown  in  Table  I  no  at- 
tention was  paid  to  the  temperature  of  the  solution,  and  the 
thiosulphate  was  run  in  until  the  liquid  turned  brown.  In 
every  case  the  solution  was  allowed  to  stand  until  there  was 
no  further  visible  change  of  color. 

TABLE  I. 


Exp. 

HgCl2  taken, 
calc'd  as 

Na2S203  in 

Volume 
at 

HgCl2  found, 
calc'd  as 

Error. 

Hg. 

excess. 

beginning. 

Hg. 

grm. 

cm8 

cm8 

grin. 

grin. 

1) 

0.0446 

46.28 

200 

0.0343 

0.0103- 

0.0354 

46.28 

400 

0.0326 

0.0028- 

3) 

0.0356 

44.97 

400 

0.0225 

0.0131- 

4) 

0.0345 

44.5 

100 

0.0308 

0.0037- 

5) 

0.0354 

44.43 

50 

0.0326 

0.0028- 

(6) 

0.0382 

22.59 

50 

0.0354 

0.0028- 

7) 

0.0375 

8.58 

50 

0.0385 

0.0010+ 

8) 

0.0371 

1.84 

50 

0.0304 

0.0067- 

9) 

0.0731 

2.28 

50 

0.0774 

0.0043+ 

(10) 

0.1486 

9.34 

50 

0.1489 

0.0003+ 

A  glance  at  the  table  shows  that  the  results  are  most 
irregular.  In  Table  II  is  seen  the  result  of  regulating  the 
temperature  and  the  length  of  standing  after  the  addition  of 
the  sodium  thiosulphate. 

*  Amer.  Chem.  Jour.,  i,  321. 


OF  THE 

UNIVERSITY 

OF 

s£*UFC 


BY  SODIUM  THIOSULPHATE. 
TABLE  II. 


331 


Exp. 

HgCl, 
taken  as 
Hg. 

Volume 
at  be- 

ginning* 

Temper- 
ature. 

Stand- 
ing. 

Na^O, 
excess. 

HgCl,  as 
Hg 
found. 

Error. 

gnn. 

cm3 

C. 

minutes. 

cm3 

grin. 

grin. 

(1) 

0.0738 

50 

36° 

40 

16.68 

0.0494 

0.0244- 

2) 
(3) 

0.0741 
0.0741 

60 
75 

70° 
70° 

15 
12 

15.42 
16.07 

0.0738 
0.0733 

0.0003- 
0.0008- 

(4) 

0.0744 

50 

70° 

10 

14.6 

0.0755 

0.0011+ 

(5) 

0.0764 

50 

72° 

7 

6.77 

0.0771 

0.0007+ 

(6) 

0.0762 

50 

75° 

10 

8.54 

0.0799 

0.0037+ 

(7) 

0.0756 

60 

73° 

15 

9.99 

0.0815 

0.0059+ 

8 

0.0774 

50 

68° 

15 

10.84 

0.0767 

0.0007- 

(9) 

0.0745 

76 

69° 

7 

6.62 

0.0805 

0.0060+ 

(10) 

0.0736 

50 

68° 

5 

15.82 

0.0714 

0.0022- 

These  results,  although  better  than  those  of  Table  I,  are 
still  very  uncertain.  On  the  supposition  that  the  change 
from  white  to  black,  which  takes  place  in  the  solution  after 
the  addition  of  an  excess  of  sodium  thiosulphate  more  or  less 
quickly  according  to  the  temperature,  was  due  to  an  increased 
amount  of  HgS  in  the  compound  2HgS  .  HgCl2,  the  next 
step  was  to  ascertain  whether  this  could  be  avoided  by 
stopping  the  addition  of  the  thiosulphate  at  the  first  indica- 
tion of  a  change  of  color  in  the  white  precipitate,  diluting 
the  solution  with  a  large  amount  of  cold  water  and  immed- 
iately throwing  it  on  the  filter.  Table  III  shows  the  result 
of  the  experiments. 

TABLE  III. 


Exp. 

HgCl2 
taken  as 
Hg. 

Volume 
at  begin- 
ing. 

Tempera- 
ture. 

Na2S203 
in 
excess. 

HgCl, 
found  as 
Hg. 

Error. 

gnn. 

cm3 

C. 

cm8 

grm. 

grm. 

(1) 

0.0749 

50 

70° 

4.15 

0.0751 

0.0002+ 

(2) 

0.0749 

50 

75° 

0.72 

0.0728 

0.0021- 

(3) 

0.0756 

50 

72° 

1.46 

0.0759 

0.0003+ 

(4) 

0.0753 

60 

70° 

2.57 

0.0750 

0.0003- 

(5) 

0.0390 

50 

70° 

3.43 

0.0396 

0.0005+ 

(6) 

0.0388 

60 

72° 

8.19 

0.0390 

0.0002+ 

(7) 

0.0380 

60 

76° 

2.03 

0.0393 

0.0013+ 

(8 

0.1494 

50 

78° 

4.99 

0.1498 

0.0004+ 

(9 
(10) 

0.1489 
0.1480 

150 
60 

78° 
70° 

4.38 
0.52 

0.1512 
0.1438 

0.0023+ 
0.0042- 

(11) 

0.1498 

50 

78° 

1.47 

0.1540 

0.0042+ 

(12) 

0.1484 

50 

71° 

2.09 

0.1517 

0.0033+ 

(13) 

0.1480 

76 

72o 

1.59 

0.1509 

0.0029+ 

332 


THE  TITRATION  OF  MERCURY 


In  the  case  of  quantities  of  mercuric  chloride  up  to  0.1 
grm.  the  results  shown  in  Table  III  are  very  satisfactory,  but 
when  larger  amounts  of  mercuric  chloride  are  used  the  errors 
again  become  prominent.  In  Table  IV,  the  effect  of  lowering 
the  temperature  to  60°  C.  and  of  increasing  the  dilution  to 
100  cm3  is  shown. 

TABLE  IV. 


Exp. 

HgCl2 
taken  as 

Volume 
at  begin- 

Temper- 
ature. 

Na2S2Os 
in 

HgCL 
found  as 

Error. 

Hg. 

ing. 

excess. 

Hg. 

grm. 

cm3 

C. 

cm3 

grm. 

grm. 

(1) 

0.0759 

100 

60° 

3.06 

0.0766 

0.0007+ 

(2) 

0.0384 

100 

60° 

2.81 

0.0387 

0.0003+ 

(3) 

0.1492 

100 

60° 

1.1 

0.1500 

0.0008+ 

w 

0.1503 

100 

60° 

1.63 

0.1506 

0.0003+ 

5 

0.1479 

100 

60° 

2.41 

0.1480 

0.0001+ 

n 

0.1489 

100 

60° 

2.12 

0.1503 

0.0014+ 

(7) 

0.2244 

100 

60° 

2.63 

0.2259 

0.0015+ 

(8 

0.1490 

100 

60° 

2.33 

0.1484 

0.0006- 

(9 
(10) 

0.0758 
0.0383 

100 
100 

60° 
60° 

2 
2.58 

0.0762 
0.0379 

0.0004+ 
0.0004- 

From  this  table  it  is  plain  that  Scherer's  process  for  the 
estimation  of  mercury  in  the  form  of  mercuric  chloride  is 
capable  of  yielding  accurate  results  if  carried  out  under 
certain  fixed  conditions.  These  conditions,  which  must  be 
closely  adhered  to,  are  as  follows:  The  solution  containing 
the  mercury  in  the  form  of  mercuric  chloride  is  placed  in  a 
liter  flask,  diluted  to  100  cm3  and  heated  to  a  temperature 
of  60°  C.  The  sodium  thiosulphate  in  ^  normal  solution 
is  run  in  from  a  burette  until  the  white  precipitate  formed 
begins  to  take  on  a  brownish  tinge.  The  solution  is  then 
diluted  with  cold  water,  some  asbestos  fiber  added  to  coagu- 
late the  precipitate  and  the  whole  is  quickly  thrown  on  the 
filter.  After  careful  washing  of  the  precipitate,  the  filtrate 
is  diluted  to  a  definite  volume,  3  grm.  of  potassium  iodide 
added  and  the  excess  thiosulphate  titrated  with  iodine  and 
starch  solution.  The  duration  of  the  process  need  not 
exceed  15  minutes.  It  is  worthy  of  note  that  there  is  no 
necessity  of  using  any  hydrochloric  acid  in  addition  to  that 


BY  SODIUM  THIOSULPHATE. 


333 


formed  in  the  reaction.  This  certainly  eliminates  one  prob- 
able source  of  error  —  the  interaction  of  hydrochloric  acid 
and  sodium  thiosulphate. 

In  dealing  with  the  estimation  of  mercury  in  the  form  of 
mercurous  nitrate  the  same  procedure  was  employed  as  in  the 
case  of  mercuric  chloride.  A  solution  of  mercurous  nitrate 
was  prepared  by  dissolving  as  much  as  possible  of  20  grm. 
of  the  salt  hi  about  200  cm3  of  water,  filtering  off  the  clear 
liquid  and  diluting  to  a  definite  volume.  The  standard  of 
the  solution  was  determined  by  precipitation  as  metallic 
mercury  by  means  of  the  electric  current.  Contrary  to  the 
statement  made  in  Sutton,  the  brown  precipitate  of  Hg2S, 
formed  as  shown  in  the  equation, 

Hg2(N08)2  +  Na2S208  =  Hg2S  +  Na2S04  +  N206, 

does  not  settle  and  leave  a  clear  supernatant  liquid,  but  the 
solution  remains  cloudy  and  it  is  impossible  to  see  any  end 
reaction.  Although  the  conditions  of  dilution,  temperature 
and  amount  of  acid  present  were  carefully  considered,  no 
arrangement  or  adjustment  of  these  conditions  was  found 

TABLE  V. 


Exp. 

Hg2(N03)a 
taken 

as  Hg. 

HNOS 
1:3. 

Volume 
at  be- 
ginning. 

Temper- 
ature. 

Na.8,0, 

in  excess. 

Hg2(NOs)2 
found 
asHg. 

Error. 

grin. 

cms 

cm* 

C. 

cm3 

grm. 

grm. 

(1) 

0.0148 

None. 

60 

60° 

4.28 

0.0129 

0.0019- 

(2) 

0.0148 

None. 

75 

60° 

4.45 

0.0117 

0.0031- 

(3) 

0.2976 

None. 

300 

60° 

12.67 

0.2760 

0.0216— 

(4) 

0.1488 

None. 

100 

40° 

2.19 

0.1386 

0.0102- 

(5) 

0.1488 

None. 

100 

50° 

6.92 

0.1378 

0.0110- 

6) 

0.1488 

None. 

200 

50° 

0.78 

0.1388 

0.0100- 

7 

0.0744 

None. 

100 

65° 

0.73 

0.0636 

0.0108- 

8 

0.0744 

1 

100 

55° 

0.49 

0.0733 

0.0011- 

9 

0.0744 

2 

100 

40° 

1.16 

0.0660 

0.0084- 

10) 

0.0744 

1 

100 

65° 

1.35 

0.0685 

0.0059- 

11) 

0.0744 

1 

100 

65° 

0.82 

0.0686 

0.0058- 

12) 

0.0744 

i 

200 

55° 

1.71 

0.0686 

0.0059- 

13 

0.0744 

100 

40° 

1.77 

0.0649 

0.0095- 

14) 

0.0744 

5* 

100 

40° 

1.34 

0.0645 

0.0099- 

15) 

0.0744 

1 

100 

45° 

1.71 

0.0654 

0.0090- 

16) 

0.0744 

10 

100 

45° 

1.55 

0.0669 

0.0075- 

(17) 

0.1488 

1 

100 

40° 

0.73 

0.1391 

0.0097- 

334  THE   TITRATION  OF  MERCURY 

under  which  satisfactory  results  could  be  obtained.  Table  V 
gives  the  result  of  experiments. 

The  errors  in  experiments  (1),  (3)  to  (6),  (10)  to  (12), 
and  (17)  are,  proportionally  to  the  amount  of  material  handled, 
practically  the  same  and  this  fact  caused  me  to  make  a  careful 
revision  of  all  standards;  but  no  mistake  could  be  found. 
The  reaction  upon  which  the  process  depends  requires  the 
formation  of  Hg2S,  but  this  mercurous  sulphide  breaks  down 
immediately  into  mercuric  sulphide  and  mercury.  The  latter 
is  probably  acted  upon  by  the  free  nitric  acid  present  to 
form  mercuric  nitrate,  which  in  turn  is  transformed  into  the 
compound  2HgS  .  Hg(NO3)2  by  the  action  of  the  thiosulphate. 
At  any  rate,  with  an  error  so  large,  whatever  its  source  may 
be,  the  process  is  plainly  impracticable. 

The  third  step  in  Scherer's  process  deals  with  the  action  of 
sodium  thiosulphate  on  mercuric  nitrate.  In  the  following 
experiments  a  solution  of  mercuric  nitrate  was  prepared 
either  by  dissolving  as  far  as  possible  20  grm.  of  mercuric 
nitrate  hi  about  200  cm3  of  cold  water,  filtering  off  the 
supernatant  liquid  and  diluting  to  a  definite  volume  (l)-(8), 
or  by  dissolving  the  salt  in  strong  nitric  acid  and  diluting 
(9)-(19).  The  standard  of  the  solution  was  obtained  by 
precipitation  as  metallic  mercury  by  means  of  the  electric  cur- 
rent. The  yellow  precipitate,  formed  according  to  Scherer's 
reaction, 

3Hg(N03)2  +  2Na2S2Os  =  2HgS  .  Hg(ST03)2  +  2Na2S04  +  2Na06, 

on  adding  the  sodium  thiosulphate  settles  much  better  than 
in  the  case  of  either  mercuric  chloride  or  mercurous  nitrate ; 
but,  as  the  supernatant  liquid  takes  on  a  permanent  yellow 
color  towards  the  end  of  the  reaction,  it  is  impossible  to  see 
when  the  thiosulphate  produces  no  further  precipitation.  On 
this  account,  therefore,  the  same  procedure  was  adopted  as 
in  the  case  of  the  mercuric  chloride  and  mercurous  nitrate, 
i.  e.,  filtration  and  titration  of  the  excess  of  sodium  thio- 
sulphate with  iodine  and  starch  solution.  The  result  of  the 
experiments  is  shown  in  the  following  table. 


BY  SODIUM  THIOSULPHATE. 
TABLE  VL 


335 


Bxp. 

Hg(N03), 
taken  as 

Volume 
at  begin- 

HN08 
1  :3 

Temper- 

NA2S203 
in 

HgfNOs), 
found  as 

Error. 

Hg. 

ing. 

excess. 

Hg. 

grm. 

cms 

cm3 

C. 

cm8 

grm. 

grm. 

(1) 

0.1167 

200 

None. 

60° 

0.46 

0.1384 

0.0217+ 

(2) 

0.1167 

100 

1 

60° 

3.63 

0.1348 

0.0181+ 

(3) 

0.1167 

100 

2 

60° 

0.23 

0.1375 

0.0208+ 

(4) 

0.1167 

100 

None. 

60° 

0.17 

0.1360 

0.0193+ 

(5) 

0.1167 

300 

None. 

21° 

0.12 

0.1232 

0.0065+ 

(6) 

0.1167 

300 

None. 

21° 

0.15 

0.1375 

0.0208+ 

(7) 

0.1167 

200 

None. 

21° 

11.8 

0.1461 

0.0294+ 

(8) 

0.1167 

200 

20 

21° 

15.97 

0.1403 

0.0236+ 

(9) 

0.1278 

200 

None. 

21° 

5.23 

0.1647 

0.0369+ 

(10) 

0.1278 

200 

None. 

21° 

1.35 

0.1653 

0.0375+ 

(11) 

0.1278 

100 

None. 

21° 

1.43 

0.1662 

0.0384+ 

(12) 

0.0752 

200 

None. 

21° 

2.09 

0.0996 

0.0244+ 

(13) 

0.0255 

100 

None. 

21° 

2.01 

0.0280 

0.0025+ 

(14) 

0.0256 

200 

5 

210 

3.44 

0.0264 

0.0009+ 

(15) 

0.0255 

200 

10 

21° 

0.85 

0.0334 

0.0079+ 

(16) 

0.0255 

300 

None. 

21° 

1.96 

0.0264 

0.0009+ 

(17) 

0.0255 

200 

None. 

21° 

1.9 

0.0287 

0.0032+ 

(18) 

0.0255 

200 

None. 

21° 

1.76 

0.0290 

0.0035+ 

(19) 

0.0639 

200 

None. 

60° 

1.53 

0.0831 

0.0192+ 

These  results  seem  to  show  the  impossibility  of  obtaining 
accurate  results  according  to  Scherer's  method  for  the  deter- 
mination of  mercuric  nitrate  by  direct  titration  with  sodium 
thiosulphate.  The  constant  plus  error  cannot  be  accounted 
for  on  the  hypothesis  that  the  nitric  acid  present  decomposes 
the  sodium  thiosulphate,  for  in  that  case  the  error  would 
lie  in  the  other  direction.  It  is  more  probable  that  the 
constitution  of  the  compound  2HgS  .  Hg(NO3)2  is  not  definite 
enough  to  make  it  the  basis  for  an  analytical  process. 


XLIH 

THE  IODOMETRIC  ESTIMATION  OF  ARSENIC 

ACID. 

BY  F.  A.  GOOCH  AND  JULIA  C.  MOKRIS.* 

THE  interaction  of  a  soluble  arseniate  and  a  soluble  iodide  in 
a  suitably  acidulated  solution  results,  as  is  well  known,  in  the 
reduction  of  the  arsenic  acid  (more  or  less  completely  according 
to  conditions  of  temperature  and  proportions  of  reagents  and 
solvents)  with  the  corresponding  liberation  of  two  atoms  of 
iodine  for  every  molecule  of  arsenic  acid  (H8O3AsO)  reduced. 
Inasmuch,  however,  as  the  reaction  of  this  process  is  reversible, 
it  is  necessary,  hi  order  that  the  reduction  may  be  complete, 
to  nullify  the  oxidizing  action  of  the  iodine  liberated. 
Theoretically  this  end  may  be  accomplished  in  either  of  two 
ways,  by  volatilizing  the  free  iodine  bodily  or  by  destroying 
the  oxidizing  power  of  the  iodine  by  converting  it  to  hydriodic 
acid.  The  former  method  was  followed  in  a  processs  devised 
for  the  estimation  of  arsenic  acid  and  elaborated  in  this 
laboratory.!  This  method,  as  originally  put  forward,  consisted 
in  adding  to  the  solution  of  the  arseniate  potassium  iodide  in 
excess  of  the  amount  theoretically  indicated,  with  10  cm3  of 
sulphuric  acid  of  half  strength,  and  so  arranging  the  dilution 
that  the  total  volume  of  the  liquid  should  be  about  100  cm3, 
boiling  until  the  volume  decreased  to  40  cm3,  bleaching  by 
the  cautious  addition  of  sulphurous  acid  the  trace  of  free 
iodine  still  held  by  the  hydriodic  acids,  diluting,  cooling, 
neutralizing  with  acid  potassium  carbonate,  and  titrating  with 
iodine,  after  adding  the  starch  indicator.  This  process, 
depending  upon  the  removal  by  volatilization  of  all  but  the 

*  From  Am.  Jour.  Sci.,  x,  151. 

t  Gooch  and  Browning,  Am.  Jour.  Sci.,  xl,  66.    Volume  I,  p.  30. 


ESTIMATION  OF  ARSENIC  ACID.  337 

last  traces  of  liberated  iodine  and  the  conversion  of  this 
minute  residue  by  sulphurous  acid,  involves  no  secondary- 
reactions  of  a  sort  likely  to  influence  the  main  effect.  It  is 
exact  and  fairly  rapid. 

The  method  of  Williamson,*  brought  forward  more  recently, 
depends  upon  the  conversion  of  the  liberated  iodine  to 
hydriodic  acid.  The  interaction  at  ordinary  temperatures  of  a 
suitably  strong  acid,  hydrochloric  or  sulphuric  acid,  upon  the 
mixture  of  the  arseniate  and  iodide  sets  free  iodine,  and  the 
liberated  iodine  is  converted  to  hydriodic  acid  by  the  action  of 
sodium  thiosulphate,  the  end  point  being  the  disappearance 
of  the  iodine  color. 

According  to  Williamson's  directions,  25  cm3  portions  of  the 
solution  of  the  arseniate  are  treated  with  potassium  iodide  and 
mixed  with  an  equal  volume  of  hydrochloric  acid  of  sp.  gr. 
1.16.  The  precaution  is  recommended  that  the  strength  of 
the  solution  of  the  arseniate  shall  not  exceed  the  decinormal 
value,  in  order  that  the  dilution  consequent  upon  titration  by 
the  thiosulphate  may  not  be  too  great  —  the  reducing  action 
brought  about  by  the  action  of  the  strong  acid  upon  the 
arseniate  and  iodide  being  reversible  upon  the  dilution  of  liquid 
with  water.  This  procedure  thus  limits  the  process  to  the 
determination  of  about  0.18  grm.  of  arsenic  acid  in  25  cm3  of 
the  solution  to  be  treated  with  an  equal  volume  of  hydrochloric 
acid  of  sp.  gr.  1.16.  Obviously,  however,  the  process  should, 
so  far  as  the  reduction  is  concerned,  be  applicable  to  larger 
amounts  of  arsenic  provided  the  strength  of  the  acid  is  kept  up 
proportionately.  It  is  essential  that  the  liquid  at  the  end  of 
the  titration  should  contain  approximately  ten  per  cent  of  its 
mass  of  absolute  hydrochloric  acid  or  about  one-third  of  its 
volume  of  the  aqueous  acid  of  sp.  gr.  1.16. 

The  arsenic  acid  is  measured  either  by  the  amount  of 
standard  thiosulphate  required  to  bleach  the  iodine  or  by  the 
amount  of  iodine  required  to  reoxidize  the  reduced  arsenious 
acid,  after  neutralizing  with  acid  potassium  carbonate.  If  the 
former  alternative  is  followed,  the  end-reaction  must  be  the 

*  Jour.  Soc.  Dyers  and  Colorists,  1896, 86-89. 
VOL.  ii.  — 22 


338 


THE  IODOMETRIC  ESTIMATION 


disappearance  of  the  yellow  color  of  the  iodine,  since  in 
solutions  so  strongly  acid  it  is  impossible  to  place  dependence 
upon  the  starch  indicator;  in  using  the  latter  alternative,  the 
starch  indicator  is,  of  course,  permissible  and  preferable. 

In  the  direct  titration  of  the  iodine  by  thiosulphate  two 
sources  of  error  present  themselves  as  possibilities ;  first,  the 
excessive  liberation  of  iodine  by  the  action  of  air  upon  the 
strongly  acidulated  iodide;  and  second,  the  liability  of  the 
thiosulphate,*  if  present  even  in  momentary  or  local  excess 
during  the  process  of  titration,  to  break  down  under  the  action 
of  strong  acid,  thus  changing  its  capacity  to  convert  iodine  to 
hydriodic  acid.  The  latter  contingency  should  be  remote  in 
proportion  to  the  caution  used  in  adding  the  thiosulphate  and 
in  keeping  the  liquid  well  stirred ;  the  former  must  of  necessity 
vary  with  the  acidity  of  the  solution  containing  the  iodide,  the 
time  of  exposure  to  atmospheric  action,  and  the  degree  of 
contact  with  the  air  incidental  to  stirring.  We  have  thought 
it  desirable,  therefore,  to  see  how  far  each  of  these  possibilities 
is  likely  to  interfere  in  the  practical  conduct  of  an  ordinary 
analysis. 


HCl 
(sp.  gr.  1.16) 
taken. 

KI 

taken. 

Total 
volume. 

Na-jSaOg  added  at 
once.    In  terms 
of  H3O3AsO. 

Na2S203  added 
after  5  minutes. 
In  terms  of 
H303AsO. 

Na2S203  added 
after  stirring 
5  minutes.     In 
terms  of 
H3O3AsO. 

cms 

grin. 

cm8 

grm. 

grin. 

grm. 

25 

2 

50 

0.0013 

. 

. 

25 

2 

75 

0.0004 

t 

25 

2 

50 

0]0035 

25 

2 

75 

0.0019 

25 

2 

50 

, 

. 

0^0042 

25 

2 

75 

. 

0.0021 

50 

2 

100 

o!ooi7 

. 

( 

60 

2 

150 

0.0004 

. 

§ 

50 

2 

100 

o!o035 

t 

50 

2 

150 

0.0019 

( 

50 

2 

100 

§ 

. 

o!o035 

50 

2 

150 

.  .  . 

0.0014 

The  effects  likely  to  result  simply  from  the  strong  acidifica- 
tion of  the  solution  containing  potassium  iodide  and  their 
variation  for  conditions  of  dilution  representing  the  beginning 
*  Norton,  Am.  Jour.  Sci.,  vii,  287.    This  volume,  p.  206. 


OF  ARSENIC  ACID. 


339 


and  the  end  of  a  titration  on  the  lines  laid  down  are  shown  in 
the  preceding  table.  The  solution  of  potassium  iodide  was 
diluted  as  indicated  before  the  addition  of  the  acid  and  the 
iodine  set  free  was  titrated  by  thiosulphate.  The  proportion- 
ate strength  of  acid  and  the  time  before  titration  are,  obviously, 
the  essential  factors.  The  absolute  amount  of  acid  present 
and  the  stirring  seem  to  make  little  difference. 

As  to  the  action  of  the  hydrochloric  acid  on  small  amounts 
of  the  thiosulphate,  we  have  the  evidence  of  the  experiments 
detailed  in  the  following  statements  in  which  1,  2,  and  5  cm3 
of  nearly  ^  thiosulphate  are  exposed  to  the  action  of  25  cm3 
hydrochloric  acid,  sp.  gr.  1.16,  without  dilution  or  diluted  with 
an  equal  volume  of  water,  were  titrated  with  nearly  ^  iodine. 
The  condition  of  acidity  when  the  volume  of  50  cm3  contains 
25  cm3  of  hydrochloric  acid,  sp.  gr.  1.16,  is  that  of  the 
beginning  of  titration  of  Williamson's  process.  In  order  that 
the  effect  of  error  due  to  such  action  upon  the  determination 
of  arsenic  acid  may  appear  immediately,  the  thiosulphate  and 
iodine  used  are  expressed  in  terms  of  that  acid. 


Error  of 

Iodine  to 

Error  of 

HCl 

Tiff 

Volume 
before 
titration. 

Na-jS-jOg  nearly 
^  in  terms 
of  H3O3AsO. 

Iodine  to 
color  with- 
out dilution, 
in  terms  of 

titration 
without 
dilution, 
in  terms  of 

color  after 
diluting 
to  75  cm8, 
in  terms  of 

titration 
after 
dilution,  in 
terms  of 

H8O3AsO. 

H3O3AsO. 

H3O3AsO. 

H808AsO. 

cms 

cm3 

cm3 

gnn. 

grm. 

grm. 

gnu. 

grm. 

25 

26 

1 

0.0071 

0.0062 

0.0009- 

0.0071 

0.0000 

25 

50 

1 

0.0071 

0.0071 

0.0000 

0.0071 

0.0000 

25* 

50 

1 

0.0071 

0.0079 

0.0008+ 

0.0079 

0.0008+ 

25 

50 

2 

0.0141 

0.0146 

0.0005+ 

0.0146 

0.0005+ 

25* 

50 

2 

0.0141 

0.0157 

0.0016+ 

0.0157 

0.0016+ 

25 

30 

5 

0.0353 

0.0336 

0.0017- 

0.0374 

0.0021+ 

25 

60 

6 

0.0353 

0.0359 

0.0006+ 

0.0359 

0.0006+ 

25* 

50 

5 

0.0353 

0.0411 

0.0058+ 

0.0411 

0.0058+ 

These  two  sources  of  error,  the  one  due  to  a  liberation  of 
iodine  and  the  other  due  to  decomposition  of  the  thiosulphate, 
would  naturally  tend  to  overcome  one  another,  but  the  com- 
pleteness of  such  neutralization  would  naturally  be  largely  a 

*  In  these  experiments  the  acid  stood  in  contact  with  the  thiosulphate 
5  minutes  before  titration. 


340 


THE  IODOMETRIC  ESTIMATION 


matter  of  chance  in  the  varying  conditions  of  actual  analysis. 
The  experiments  of  the  following  table,  however,  in  which 
^  thiosulphate,  to  the  amount  of  1,  2,  and  5  cm8,  was  added 
to  the  liquid,  50  cm3  and  75  cm3,  containing  25  cm3  acid,  and 
titrated  with  iodine  at  once,  and  after  five  minutes,  were 
made  to  test  the  matter  for  the  conditions  of  dilution  at  the 
beginning  and  at  the  end  of  a  titration. 


HCl 
(sp.  gr. 
1.16). 

KI. 

Volume. 

Na2S2Oa  nearly 
JC.  in  terms 
of  H303AsO. 

Iodine  in 
terms  of 
H3OsAsO, 
at  once. 

Iodine  in 
terms  of 
H3O3AsO, 
after  5  min. 

Error 
in  terms  of 
H3OsAsO. 

cm3 

grm. 

cm8 

cm» 

grm. 

grm. 

grm, 

grm. 

25 

2 

50 

1 

0.0071 

0.0057 

0.0014- 

25 

2 

76 

1 

0.0071 

0.0071 

. 

0.0000 

25 

2 

50 

2 

0.0141 

00131 

t 

0.0010- 

25 

2 

76 

2 

0.0141 

00143 

.       . 

0.0002+ 

25 

2 

50 

5 

0.0353 

00332 

f 

0.0021- 

25 

2 

75 

5 

0.0353 

00357 

. 

0.0004+ 

25 

2 

50 

1 

0.0071 

0.0028 

0.0048- 

25 

2 

76 

1 

0.0071 

0.0067 

0.0004- 

25 

2 

50 

2 

0.0141 

t 

0.0116 

0.0025- 

25 

2 

75 

2 

0.0141 

t 

0.0139 

0.0002- 

25 

2 

60 

6 

0.0353 

0.0312 

0.0041- 

25 

2 

75 

6 

0.0353 

• 

0.0361 

0.0008+ 

It  is  clear  that  under  the  conditions  covered  by  the  experi- 
ments of  the  two  preceding  tables  the  decomposition  of  the 
thiosulphate  is  likely  to  occur  in  greater  or  less  degree,  and 
that  when  the  acid  of  sp.  gr.  1.16  is  not  much  diluted,  the 
products  of  decomposition  are  not  oxidized  by  the  iodine 
completely.  The  latter  observation  is  quite  hi  harmony 
with  the  fact  that  sulphur  dioxide  bleaches  iodine  in  strong 
hydrochloric  acid  only  slowly  and  incompletely.  In  such 
cases  dilution  favors  further  action  of  the  iodine,  but  results 
obtained  by  titration  with  iodine  in  the  acid  solution  diluted 
with  an  equal  amount  of  water  are  unmodified  by  further 
dilution. 

In  the  following  tables  are  recorded  actual  determinations 
of  arsenic  according  to  Williamson's  process.  To  each  25 
cm3  of  the  arseniate  were  added  1,  2,  or  3  grms.  of  potassium 
iodide  and  25  cm3  hydrochloric  acid,  sp.  gr.  1.16.  The  iodine 


OF  ARSENIC  ACID. 


341 


was  bleached  by  nearly  decinonnal  thiosulphate  without 
addition  of  the  starch  indicator,  which  loses  all  delicacy  in 
the  presence  of  strong  acid.  The  time  occupied  by  each 
titration  was  about  five  minutes.  The  standards  of  the 
arseniate  were  determined  by  the  vaporization  process,*  the 
purity  of  reagents  employed  in  that  process  having  been 
proved  by  trying  the  process  in  the  estimation  of  a  solution 
of  arsenic  acid  made  by  oxidizing  pure  decinormal  arsenious 
acid  by  iodine. 


Volume  at 

Volume 

H2KAs04 

HCL 

KI. 

beginning 
of 

at 

end  of 

m 
terms  of 

Hs03AsO 
found. 

Error. 

titration. 

titration. 

H3O3AsO. 

cm3. 

grin. 

cm8. 

cms. 

grin. 

grm. 

grm. 

25 

2 

50 

51 

0.0062 

0.0085 

0.0023+ 

25 

2 

50 

52 

0.0125 

0.0156 

0.0031+ 

25 

2 

50 

55 

0.0312 

0.0350 

0.0038+ 

25 

2 

50 

55 

0.0624 

0.0666 

0.0042+ 

25 

2 

50 

73 

0.1559 

0.1588 

0.0029+ 

25 

2 

60 

73 

0.1559 

0.1587 

0.0028+ 

25 

2 

60 

73 

0.1559 

0.1591 

0.0032+ 

25 

2 

50 

73 

0.1559 

0.1595 

0.0036+ 

25 

3 

50 

73 

0.1559 

0.1595 

0.0036+ 

25 

1 

50 

73 

0.1559 

0.1581 

0.0022+ 

25 

2 

50 

73 

0.1559 

0.1581 

0.0022+ 

25 

2 

50 

73 

0.1559 

0.1588 

0.0029+ 

The  range  of  error  in  these  results  is  from  -f-  0.0022  grm. 
to  +  0.0042  grm.  with  a  mean  of  +  0.0031  grm.  —  not  very 
different  from  what  might  be  expected  from  the  effect  of  the 
interaction  of  the  strong  hydrochloric  acid  and  the  iodide 
alone.  The  counter-effect  due  to  the  decomposition  of  the 
thiosulphate  is  not  large,  yet  it  is  probably  real,  as  will  appear 
in  the  sequel.  In  the  following  series  of  determinations, 
made  with  new  solutions  and  new  standards  throughout,  the 
arsenic  acid  was  determined,  first,  by  titrating  the  iodine  set 
free  by  25  cm3  of  hydrochloric  acid,  sp.  gr.  1.16  and  3  grms. 
potassium  iodide,  the  solution  having  a  total  volume  of  50  cm8 
at  beginning  and  of  75  cm3  at  the  end  of  titration  and, 
secondly,  the  arsenious  acid  produced  in  the  first  reaction 
was  titrated,  after  being  neutralized  with  acid  potassium  car- 
bonate by  iodine  in  the  presence  of  the  starch  indicator. 
*  Gooch  and  Browning,  loc.  cit. 


342 


THE  IODOMETRIC  ESTIMATION 


H2KAs04 

taken,  in 
terms  of 
H3O3AsO. 

H303A80 
found  by  the 
thioeulphate. 

Error. 

H303A80 

found  by 
titration  of 
HSO3A8  with 
iodine. 

Error. 

grm. 

gnn. 

grm. 

grm. 

grm. 

0.1767 

0.1798 

0.0031+ 

0.1776 

0.0009+ 

0.1767 

0.1798 

0.0031+ 

0.1777 

0.0010+ 

0.1767 

0.1795 

0.0028+ 

0.1785 

0.0018+ 

0.1767 

0.1793 

0.0026+ 

0.1785 

0.0018+ 

0.1767 

0.1794 

0.0027+ 

0.1780 

0.0013+ 

0.1767 

0.1798 

0.0031+ 

0.1785 

0.0018+ 

The  average  error  of  the  first  operation  is  0.0029  grm.,  not 
far  from  that  of  the  previous  series ;  the  error  of  the  second 
operation,  the  titration  of  the  arsenious  acid,  amounts  on  the 
average  to  0.0014  grm.  In  the  second  operation  the  error 
due  to  over-use  of  the  thiosulphate  by  iodine  set  free  outside 
the  main  reaction  is  obviously  eliminated.  The  tetrathionate 
present  after  neutralization  with  acid  potassium  carbonate  is 
unaffected  by  iodine,  as  we  have  found  by  titrating  25  cm3  ^ 
iodine  mixed  with  25  cm3  hydrochloric  acid,  sp.  gr.  1.16,  by 
the  thiosulphate,  neutralizing  with  acid  potassium  carbonate,* 
adding  starch  and  getting  the  starch  blue  with  a  single  drop 
of  -^  iodine.  The  average  error  of  this  process,  therefore, 
0.0014,  is  probably  due  to  the  products  of  decomposition  of 
the  thiosulphate  in  the  first  operation. 

From  the  foregoing  experiments  it  is  clear  that  an  arbitrary 
correction  of  about  0.0030  grm.  must  be  deducted  from  the 
indications  of  Williamson's  process  of  direct  titration  by 
thiosulphate,  made  with  the  greatest  care  under  the  conditions 
mentioned ;  and  that  a  correction  varying  from  one-half  that 
amount  (0.0015  grm.)  to  nothing  (according  to  the  amount 
of  arsenious  acid  present)  when  the  determination  is  made 
by  iodine  after  neutralization  with  acid  potassium  carbonate. 

*  It  is  worthy  of  note,  that,  as  we  have  found  by  experience,  it  is  not 
possible  to  substitute  an  alkaline  hydroxide  for  the  carbonate  in  the  early 
stages  of  the  process  of  neutralization,  on  account  of  the  decomposing  effect 
of  the  former  reagent  upon  the  tetrathionate.  This  effect  is  in  proportion  to 
the  heating  of  the  solution,  but  is  never  wholly  absent  even  when  ice  is 
intermixed  with  the  liquid  and  the  greatest  care  taken  to  prevent  a  rise  of 
temperature. 


OF  ARSENIC  ACID. 


343 


After  making  these  arbitrary  corrections  in  the  results  of  the 
preceding  table,  the  individual  variations  fall  within  reason- 
able limits. 

On  the  other  hand,  the  vaporization  process,  in  which 
the  arseniate  is  reduced  by  boiling  with  sulphuric  acid  and 
potassium  iodide  in  the  manner  described,*  gives  indications 
reasonably  regular  and  accurate  without  the  application  of 
an  arbitrary  correction.  This  process,  moreover,  may  be 
shortened  by  restricting  the  volume  at  which  heating  begins 
so  that  the  boiling  need  not  be  extended  beyond  five  or  six 
minutes.  According  to  this  slight  modification,  the  solution 
of  the  arseniate  is  heated  in  an  Erlenmeyer  flask  with  potas- 
sium iodide  to  an  amount  about  0.5  grm.  in  excess  of  the 
amount  theoretically  required  and  10  cm3  of  sulphuric  acid 
of  half  strength  in  a  total  volume  of  between  50  cm3  and 
75  cm3.  The  liquid  is  boiled  till  the  iodine  vapors  are  no 
longer  visible  in  the  flask  above  the  liquid,  the  iodine  color 
in  the  still  hot  liquid  is  bleached  by  the  cautious  addition  of 
sulphurous  acid,  the  whole  is  diluted  with  cold  water,  and 
cooled  quickly.  The  solution  is  nearly  neutralized  with 
potassium  hydroxide  and  the  neutralization  is  completed  with 
acid  potassium  carbonate.  The  reduced  acid  is  titrated  with 
iodine  after  adding  the  starch  indicator.  By  this  procedure 
the  results  of  the  following  table  were  obtained. 


Volume. 

H808ABO 
taken. 

H803AsO 
found. 

Error. 

cm8. 

grm. 

grm. 

grm. 

35 

0.1559 

0.1559 

0.0000 

35 

0.1559 

0.1560 

0.0001+ 

40 

0.1559 

0.1559 

0.0000 

65 

0.1559 

0.1559 

0.0000 

50 

0.2495 

0.2499 

0.0004+ 

50 

0.2557 

0.2449 

0.0008- 

60 

0.3119 

0.3117 

0.0002- 

60 

0.3119 

0.3120 

0.0001+ 

75 

0.3119 

0.3124 

0.0005+ 

75 

0.3119 

0.3132 

0.0013+ 

75 

0.3119 

0.3121 

0.0002+ 

75 

0.3119 

0.3115 

0.0004- 

75 

0.3119 

0.3124 

0.0005+ 

*  Loc.  cit 


XLIV 

ON  THE  QUALITATIVE  SEPARATION  OF  NICKEL 
FROM  COBALT  BY  THE  ACTION  OF  AMMO- 
NIUM HYDROXIDE  ON  THE  FERRICYANIDES. 

BY  PHILIP  E.  BROWNING  AND  JOHN  B.  HARTWELL* 

SOME  years  ago  F.  W.  Clarke  f  suggested  a  method  for  the 
separation  of  nickel  from  cobalt  depending  upon  the  solvent 
action  of  ammonium  hydroxide  upon  the  precipitated  fern- 
cyanides.  The  method  may  best  be  described  by  quoting 
from  the  original  article :  "  To  the  slightly  acid  solution 
containing  the  two  metals,  I  first  add  an  excess  of  ammonium 
chloride.  This  causes  the  cobalt  precipitate,  which  otherwise 
would  run  through  the  filter,  to  fall  in  a  denser  state,  and 
also  of  a  much  darker  color,  often  nearly  black.  I  then  add 
the  potassium  ferricyanide  until  the  precipitation  is  complete, 
and  afterwards  agitate  strongly  with  a  considerable  excess  of 
ammonia.  Upon  filtering,  all  the  cobalt  remains  upon  the 
filter,  being  recognized  by  the  characteristic  color  of  the 
precipitate,  and  the  nickel  is  readily  detected  in  the  filtrate, 
by  means  of  ammonium  sulphide.  If,  upon  filtering,  the 
portion  at  first  running  through  is  turbid,  it  may  be  disre- 
garded, or  returned  to  the  filter,  that  which  filters  through 
subsequently  being  almost  invariably  clear." 

In  making  a  study  of  this  method  we  found  two  serious 
objections;  first,  the  practical  impossibility  of  obtaining  a 
good  filtration  from  the  cobalt  ferricyanide,  even  in  the  pres- 
ence of  the  ammonium  chloride,  and,  second,  the  large  amount 
of  sulphur  thrown  down  when  ammonium  sulphide  was 

*  From  Am.  Jour.  Sci.,  x,  316. 
t  Am.  Jour.  Sci.,  xlviii,  67. 


SEPARATION  OF  NICKEL  FROM  COBALT.         345 

added  to  the  filtrate  containing  the  nickel  with  the  excess 
of  ferricyanide. 

Our  first  attempt  was  to  secure,  if  possible,  a  complete 
separation  of  the  precipitated  cobalt  ferricyanide  and  the 
dissolved  nickel  by  filtration.  This  we  were  able  to  accom- 
plish by  the  addition  of  a  small  amount  of  a  solution  of  an 
aluminum  salt  to  the  original  solution  which  held  back  the 
cobalt,  and,  as  experiment  showed,  allowed  the  complete 
solvent  action  of  the  ammonium  hydroxide  upon  the  nickel 
salt.  Amounts  of  nickel  as  small  as  0.0001  grm.  were 
detected,  when  mixed  with  the  aluminum  salt,  by  precipi- 
tating as  ferricyanide,  extracting  with  ammonium  hydroxide, 
and  testing  in  the  manner  to  be  described. 

On  turning  our  attention  to  a  possible  improvement  hi  the 
method  for  the  detection  of  the  nickel,  a  reaction  first  dis- 
cussed by  Allen  *  was  applied.  When  the  ammoniacal  solution 
of  the  nickel  ferricyanide  was  treated  with  strong  sodium  or 
potassium  hydroxide  solution,  in  the  presence  of  an  excess 
of  potassium  ferricyanide,  a  black  flocky  precipitate  formed 
which  gave  no  test  for  ferro-  or  ferricyanide,  and  gave  every 
indication  of  being  nickelic  hydroxide.  This  reaction  we 
found  to  afford  us  a  most  delicate  test  for  nickel. 

The  method  as  modified  by  us  may  be  described  as  follows : 
Dissolve  not  more  than  0.1  grm.  of  the  salts  of  the  two  ele- 
ments hi  about  5  cm3  of  water,  add  a  few  drops  of  a  saturated 
solution  of  alum,  destroy  any  free  mineral  acid  by  neutraliz- 
ing with  ammonium  hydroxide,  and  make  faintly  acid  with 
acetic  acid.  To  this  solution  add  about  0.5  grm.  of  potassium 
ferricyanide  and  agitate  to  effect  the  solution  of  the  ferri- 
cyanide and  the  complete  precipitation  of  the  nickel  and  cobalt 
salts.  Then  add  about  5  cm3  of  strong  ammonium  hydroxide 
and  filter.  To  the  filtrate,  which  should  have  no  reddish 
color,  add  a  piece  of  sodium  or  potassium  hydroxide  about 
the  size  of  a  pea  and  boil.  The  appearance  of  a  black  pre- 
cipitate, showing  first  as  a  dark  coloration  in  case  of  very 
small  amounts,  indicates  nickel. 

*  Chein.  News,  xxxiii,  290. 


346 


SEPARATION  OF  NICKEL  FROM  COBALT. 


The  tables  following  give  a  record  of  the  experimental 
results.  With  the  precautions  indicated,  this  method  may 
be  applied  very  satisfactorily. 


I. 

Exp. 

CoSOJHjjO. 

NiS04. 
7H80. 

KA1(S04)2 
saturated 
solution. 

K3FeC8N6 

NH4OH 

concen- 
trated. 

NaOH  solid. 

Result. 

griu. 

gnn. 

cm» 

grin. 

cms 

(1) 

.    .    . 

0.0100 

2 

0.5 

5 

About  size 

Heavy  ppt. 

of  a  pea. 

(2) 

.    .    , 

0.0050 

2 

0.5 

5 

(i 

Heavy  ppt. 

(3) 

•    •    . 

0.0010 

2 

0.5 

5 

« 

Heavy  ppt. 

(4) 

0.0003 

2 

0.5 

5 

<t 

Distinct. 

(5) 

.    .    . 

0.0001 

2 

0.5 

5 

a 

Plain. 

II. 

(1) 

0.10 

.    .    . 

2 

0.5 

5 

About  size 

None. 

of  a  pea. 

(2) 

0.10 

0.0100 

2 

0.5 

5 

Heavy. 

(3) 

0.10 

0.0050 

2 

0.5 

5 

Distinct. 

(4) 

0.10 

0.0030 

2 

0.5 

5 

Very  faint. 

(5) 

0.05 

2 

0.5 

5 

None. 

(6) 

0.05 

0.0100 

2 

0.5 

5 

Heavy. 

(7) 

0.05 

0.0050 

2 

0.5 

5 

Distinct. 

(8) 

0.05 

0.0030 

2 

0.5 

5 

Plain. 

(9) 

0.05 

0.0010* 

2 

0.5 

5 

Faint. 

*  Equivalent  to  0.0002  of  the  metal. 


XLV 

THE  VOLUMETRIC  ESTIMATION  OF  COPPER  AS 
THE  OXALATE,  WITH  SEPARATION  FROM 
CADMIUM,  ARSENIC,  TIN,  IRON,  AND  ZINC. 

BY  CHARLES  A.  PETERS* 

IT  is  a  well  known  fact  that  copper  oxalate  is  insoluble  in 
water  and  scarcely  attacked  by  moderate  amounts  of  dilute 
nitric  acid.f  Upon  this  fact  Bornemann  ^  has  recently 
based  a  method  for  the  separation  of  copper  from  cadmium 
by  precipitating  copper  as  the  oxalate  in  the  presence  of 
nitric  acid,  filtering  hot,  and  estimating  the  copper  after 
ignition,  by  any  of  the  well  known  gravimetric  methods.  Six 
to  ten  grams  of  copper,  as  the  oxide,  were  used  for  a  single 
determination,  and  the  errors  were  large.  Bornemann  does 
not  recommend  this  process  as  an  accurate  analytical  method. 
Classen  §  describes  a  method  for  the  separation  of  metals  as 
oxalates  by  adding  to  the  solution  of  the  salt  of  the  metals 
a  dilute  solution  of  the  potassium  oxalate  (1  :  6)  and  con- 
centrated acetic  acid  to  80  per  cent  of  the  total  volume. 
Regarding  copper  salts  in  particular,  Classen  states  that 
precipitation  takes  place  only  in  dilute  solution  and  then  not 
completely. 

It  has  been  the  experience  of  the  writer,  that  the  precipi- 
tation of  copper  oxalate  from  solutions  containing  at  least 
0.0128  grm.  of  the  oxide  and  saturated  with  the  oxalic  acid  is 
practically  complete.  The  filtrate  in  such  cases  gives  no 
blue  color  with  ammonia,  looking  down  on  a  column  of  liquid 

*  From  Am.  Jour.  Sci.,  x,  359. 

t  Storer,  Dictionary  of  Chemical  Solubilities,  p.  463. 

J  Chem.  Ztg.,  xxiii,  565.  §  Ber.  Dtsch.  chem.  Ges.,  x,  6, 1316. 


348 


VOLUMETRIC  ESTIMATION  OF  COPPER. 


in  a  test-tube,  and  only  a  faint  brown  color  is  developed 
when  the  nitrate  is  neutralized,  made  acid  with  acetic  acid, 
and  tested  with  potassium  ferrocyanide.  It  is  the  object  of 
this  paper  to  show  that  moderate  amounts  of  copper  may  be 
determined  quantitatively  as  the  oxalate  by  precipitation  with 
oxalic  acid  and  titration  of  the  precipitate  by  potassium 
permanganate,  and  also  to  show  that  moderate  amounts  of 
copper  may  be  separated  from  other  metals  in  the  presence 
of  nitric  acid,  by  the  addition  of  considerable  amounts  of 
oxalic  acid. 

Before  attempting  the  quantitative  separation  of  copper 
from  solution  by  the  addition  of  oxalic  acid  a  few  qualitative 
experiments  upon  the  precipitation  of  varying  amounts  of 
copper  sulphate  by  varying  amounts  of  oxalic  acid  were  tried 
at  different  dilutions.  In  all  the  experiments  the  mixtures 
stood  16-20  hours,  and  were  filtered  from  2  to  4  times 
through  four  filters  folded  together,  and  the  nitrates  were 
tested  both  with  ammonia  and  with  potassium  ferrocyanide. 
In  cases  where  the  filtrate  gave  no  blue  color  with  ammonia 

TABLE  I. 

Dilution  50  cm8. 


Ezp. 

CuO 

taken  as 
CuSO4. 

Oxalic  acid  added  in  solution. 

Oxalic  acid  added  in  crystalline 
form. 

Filtrate 
treated 
with  NH4OH. 

Filtrate 
treated  with 

K4FeC6N6. 

Filtrated 
with 
NH4OH. 

Filtrate 
treated  with 
E^FeCeNa. 

2.0  grm. 
oxalic 
acid 
present. 

1.0  grm. 
oxalic 
acid 
present. 

0.5  grm. 
oxalic 
acid 
present. 

• 

grm. 

r  o.ois 

0.031 
0.051 
0.064 
0.018 
0.031 
0.051 
0.064 
0.094 
'  0.018 
0.031 
0.051 
0.064 
0.094 

Blue  color. 
Trace  '« 

Abundant  ppt. 
Abundant  ppt. 
Evident  ppt. 
Trace  ppt. 
Abundant  ppt. 
Abundant  ppt. 
Abundant  ppt. 
Evident  ppt. 
Trace  ppt. 
Abundant  ppt. 
Abundant  ppt. 
Abundant  ppt. 
Evident  ppt. 
Trace  ppt 

Blue  color. 
Trace  " 

Abundant  ppt. 
Evident  ppt. 
Evident  ppt. 
Trace  ppt. 
Abundant  ppt. 
Evident  ppt. 
Evident  ppt. 
Evident  ppt. 
Trace  ppt. 
Abundant  ppt. 
Abundant  ppt. 
Evident  ppt. 
Trace  ppt. 
Trade  ppt 

Blue  color. 
Blue  color. 
Trace  " 

Blue  color. 
Trace  " 

Blue  color. 
Blue  color. 
Trace  " 

Blue  color. 
Trace  " 

VOLUMETRIC  ESTIMATION  OF  COPPER. 


349 


and  only  a  slight  precipitate  with  ferrocyanide  the  precipita- 
tion was  considered  practically  complete  and  the  conditions 
were  regarded  suitable  for  the  trial  of  the  method  quantita- 
tively. In  the  following  table  is  recorded  the  work  upon  the 
precipitation  of  copper  sulphate  by  0.5  grm.,  1.0  gnn.,  and  2.0 
grm.  of  oxalic  acid  in  50  cm3  of  solution. 

It  will  be  seen  readily  by  comparison  of  the  right  and  left 
hand  sides  of  the  table  above  that  somewhat  smaller  amounts 
of  copper  may  be  precipitated  completely  by  the  addition  of 
crystallized  oxalic  acid  than  by  the  same  amount  of  oxalic 
acid  already  in  solution.  Thus,  when  dissolved  oxalic  acid 
is  added  to  the  solution  of  50  cm8,  amounts  of  copper  sulphate 
less  than  0.040-0.050  grm.  are  not  precipitated  completely, 
while  under  conditions  otherwise  the  same  excepting  that  the 
oxalic  acid  is  added  in  crystalline  form,  the  precipitation  of 
amounts  as  small  as  0.030  grm.  is  practically  complete.  The 
amount  of  oxalic  acid  in  solution  necessary  for  the  complete 

TABLE  H. 


A. 

CuO 

taken  as 
CuSO4. 

Oxalic 
acid 
added  in 
solution. 

Volume 
at  pre- 
cipita- 
tion. 

Filtrate  treated  with 
NH4OH. 

Filtrate  treated 
with  K4FeC6N8. 

gnn. 
0.031 
0.031 
0.031 
0.031 
0031 

gnn. 
0.5 
1.0 
2.0 
3.0 
35 

cm8 

50 
50 
50 
60 
50 

Blue  color. 
Blue  color. 
Trace  blue  color. 
Slight  trace  blue  color. 

Abundant  ppt. 
Abundant  ppt. 
Abundant  ppt. 
Abundant  ppt. 
Evident  ppt. 

0.0128 
0.0064 

5.0 
5.0 

50 
50 

No  blue. 
Blue  color. 

Trace  ppt. 
Abundant  ppt. 

B. 

0.0064 
0.0064 
0.0064 
00064 

0.5 
0.5 
0.5 
05 

20 
15 
10 
5 

Faint  blue. 
Faint  blue. 
Faint  blue. 

Abundant  ppt. 
Abundant  ppt. 
Faint  ppt. 
Trace  ppt. 

00003 

05 

5* 

00003 

01 

If 

*  Precipitate  redissolved. 


t  Precipitate  remained. 


350  VOLUMETRIC  ESTIMATION  OF  COPPER. 

precipitation  (after  16  to  20  hours)  of  this  minimum  amount 
of  copper,  0.031  grm.  of  copper  oxide  taken  as  the  sulphate, 
appears,  as  shown  in  Table  II,  A,  which  precedes,  to  be  about 
3.5  grm.  in  50.0  cm3.  If  the  amount  of  oxalic  acid  is  in- 
creased to  5  grm.,  making  the  solution  saturated  for  that 
substance,  using  the  same  volume  of  liquid,  the  minimum 
amount  completely  precipitable  is  reduced  to  0.0128  grm.  but 
not  to  one-half  that  amount. 

It  appears  from  the  experiments  of  Table  II,  B,  that  the 
volume  of  liquid  in  which  precipitation  takes  place  influences 
the  complete  precipitation  of  the  copper  oxalate.  Thus  the 
precipitation  of  0.0064  grm.  of  copper  oxide  taken  as  the 
sulphate  by  0.5  grm.  of  oxalic  acid  is  complete  in  5  cm3  of 
liquid.  The  precipitate  which  falls  from  0.0003  grm.  of  the 
oxide  taken  as  the  sulphate  dissolves  in  5  cm3  of  liquid,  but 
remains  visible  in  1  cm3. 

As  a  result  of  the  preluninary  experiments,  it  may  be  said 
that  the  presence  of  a  certain  minimum  amount  of  copper, 
varying  with  the  conditions,  is  essential  to  complete  precipi- 
tation. Thus,  at  a  dilution  of  50  cm3  a  saturated  solution  of 
oxalic  acid  will  precipitate  with  practical  completeness  copper 
taken  as  the  sulphate  in  amounts  exceeding  the  equivalent 
of  0.0128  grm.  of  copper  oxide;  2.0  grm.  of  oxalic  acid 
will  precipitate  almost  completely  for  the  same  volume  of 
solution  the  equivalent  of  0.03  grm.  of  copper  oxide;  and 
1.0  grm.  or  0.5  grm.  of  oxalic  acid  will  precipitate  the  equiva- 
lent of  0.064  grm.  of  the  oxide. 

In  the  quantitative  separation  of  copper  as  the  oxalate 
the  method  of  treatment  was  in  general  as  follows.  Copper 
sulphate  hi  50  cm3  of  water  was  thrown  down  by  the  addition 
of  dry  oxalic  acid  to  the  hot  solution,  and,  after  standing 
over  night,  the  precipitate  was  filtered  on  asbestos  and  washed 
two  or  three  times  with  small  amounts  of  cold  water.  The 
precipitate,  still  in  the  crucible,  was  returned  to  the  beaker 
in  which  precipitation  took  place,  5  or  10  cm3  of  dilute  sul- 
phuric acid  (1  : 1)  were  then  added  together  with  a  convenient 
amount  of  water,  and,  after  heating  the  liquid  to  boiling,  the 


VOLUMETRIC  ESTIMATION  OF  COPPER. 


351 


oxalic  acid  was  titrated  with  permanganate,  the  oxalate  of 
copper  dissolving  readily  as  fast  as  the  excess  of  oxalic  acid 
is  removed  by  the  permanganate.  The  precipitate  may  also 
be  dissolved  in  10  cm3  of  strong  hydrochloric  acid,*  and,  after 
adding  0.5  grm.  manganous  chloride,  titrated  at  30°-50°. 
Experiments  (4)  and  (5)  were  conducted  after  this  manner. 
In  Table  III,  A,  which  follows,  are  recorded  results  of  the 
quantitative  tests  of  the  method. 

TABLE  III. 


Exp. 

CuO 
taken  as 
CuSO4. 

Oxalic 
acid. 

Volume 
at  precipi- 
tation. 

CuO 
found. 

Error. 

A. 

grm. 

grm. 

cm» 

grm. 

grm. 

(1) 

0.0372 

0.15 

100 

0.0286 

0.0086- 

(2) 

0.1860 

0.50 

125 

0.1831 

0.0029- 

(3) 

0.0398 

0.50 

50 

0.0376 

0.0022- 

(4) 

0.1860 

1.0 

150 

0.1834 

0.0026- 

(5) 

0.1860 

0.5 

50 

0.1864 

0.0004+ 

(6) 

0.1860 

0.5 

60 

0.1866 

0.0006+ 

(7) 

0.1860 

0.5 

50 

0.1866 

0.0006+ 

(8) 

0.1860 

1.0 

50 

0.1866 

0.0006+ 

(9) 

0.0398 

1.0 

50 

0.0391 

0.0007- 

In  experiments  (l)-(4),  deficiencies  are  found  in  the 
amounts  of  oxalate  precipitated  at  different  degrees  of  dilution 
and  by  different  amounts  of  the  precipitant  which  are  in 
agreement  with  the  results  obtained  in  the  preliminary  work ; 
the  results  of  experiments  (5)-(9),  in  which  0.5  grm.  and  1.0 
grm.  of  oxalic  acid  act  in  a  total  volume  of  50  cm3,  show 
the  precipitation  to  be  essentially  complete  under  these 
conditions. 

To  study  the  insolubility  of  the  copper  oxalate  in  nitric 
acid  the  experiments  in  Section  B  of  the  table  were  made. 

In  experiments  (10)-(13)  amounts  of  oxalic  acid  varying 
from  0.5  grm.  to  3.0  grm.  appear  to  precipitate  the  copper 
completely  in  the  presence  of  5  cm3  of  strong  nitric  acid.  In 
experiment  (14)  the  amount  of  oxalic  acid  used  was  not  suffi- 


Gooch  and  Peters,  Am.  Jour.  Sci.,  vii,  461.    This  volume,  p.  222. 


352 


VOLUMETRIC  ESTIMATION  OF  COPPER. 


cient  to  throw  down  all  the  copper  in  the  presence  of  10  cm3 
of  nitric  acid,  but  the  copper  does  come  down  completely  in 
the  presence  of  the  large  amount  of  the  nitric  acid  upon  the 
addition  of  more  oxalic  acid,  as  seen  in  experiments  (15)  and 
(16).  In  experiments  (17)  and  (18)  with  a  larger  volume  of 
water  and  a  larger  absolute  amount,  though  approximately 
the  same  percentage,  of  nitric  acid  present  as  in  experiments 
(10)-(13),  there  is  a  slight  loss  of  copper;  but  hi  experiments 
(21)  and  (22)  when  the  amount  of  nitric  acid  is  reduced  to 
5  cm3  in  the  larger  total  volume  the  results  are  normal. 
Experiments  (19)  and  (20)  show  the  increased  loss  when  still 
larger  amounts  of  nitric  are  present.  These  facts  would  make 
it  seem  best  to  limit  the  absolute  amount  of  nitric  in  solution 
to  about  5  cm3. 

TABLE   III  (continued). 


CuO 

HNOg 

Volume 

Ezp. 

taken  as 
CuS04. 

acid. 

<!y? 

at  precipi- 
tation. 

CuO  found. 

Error. 

B. 

grm. 

gnn. 

cm3 

cm» 

gnn. 

grm. 

(10) 

0.1860 

0.5 

5.0 

55 

0.1859 

0.0001- 

(11) 

0.1860 

0.5 

5.0 

65 

0.1860 

0.0000 

(12) 

0.1990 

2.0 

5.0 

66 

0.1989 

0.0001- 

(13) 

0.1990 

3.0 

6.0 

55 

0.1990 

0.0000 

(14) 

0.1990 

2.0 

10.0 

60 

0.1971 

0.0019- 

(15) 

0.1990 

3.0 

10.0 

60 

0.1987 

0.0003- 

16 

0.1990 

3.0 

10.0 

60 

0.1985 

0.0005- 

(17) 

0.1990 

5.0 

12.0 

130 

0.1977 

0.0013- 

(18) 

0.1990 

5.0 

12.0 

130 

0.1975 

0.0015- 

(19) 

0.1990 

5.0 

25.0 

130 

0.1837 

0.0153- 

(20) 

0.1990 

5.0 

25.0 

130 

0.1831 

0.0159— 

(21) 

0.1990 

6.0 

5.0 

130 

0.1983 

0.0007- 

(22) 

0.1990 

6.0 

5.0 

130 

0.1988 

0.0002- 

C. 

(23) 

0.1990 

2.5 

5.0* 

65 

0.1961 

0.0029- 

(24) 

0.1990 

2.0 

5.0 

65 

0.1971 

0.0019- 

One  observation  may  well  be  noted  here :  namely,  that  while 
one-half  gram  oxalic  acid  is  all  that  is  needed  for  the  complete 

*  About  9  grm.  of  ammonium  nitrate  present  in  addition  to  the  5  cm3  of 
nitric  acid. 


VOLUMETRIC  ESTIMATION  OF  COPPER. 


353 


precipitation  of  the  copper  in  the  presence  of  5  cm3  strong 
nitric  acid,  still  the  oxalic  acid  may  be  added  up  to  the  point 
of  saturation  of  the  solution.  More  than  this  causes  difficulty 
owing  to  the  fact  that  a  large  amount  of  water  is  necessary  to 
wash  the  precipitated  oxalate.  About  2.0  grin,  of  oxalic  acid 
to  50  cm3  of  water  is  a  convenient  proportion. 

In  experiments  (23)  and  (24),  5  cm8  of  nitric  acid  were 
neutralized  with  ammonium  hydroxide  before  adding  the  5  cm3 
strong  nitric  acid  in  excess.  The  results  show  the  solubility 
of  copper  oxalate  in  ammonium  nitrate  and  exclude  the 
possibility  of  such  a  procedure  in  this  work. 

TABLE  HI  (continued). 


Exp. 

CuO 

taken  as 
CuS04. 

Oxalic 
acid. 

HN03 
(sp.  gr. 
1.40). 

Volume 
at  precip- 
itation. 

CuO 

found. 

Error. 

Details 
of  filtration. 

D. 

grin* 

gnu* 

cm3 

cms 

gnu. 

grin. 

(23) 

0.1990 

2.0 

50 

0.1984 

0.0006- 

Filtered  hot 
immediately. 

(24) 

0.2030 

2.0 

50 

0.2025 

0.0005- 

Filtered  hot 
immediately. 

t  Filtered  after 

(25) 

0.1990 

1.0 

.  . 

50 

0.1990 

0.0000 

?  cooling  ;  stood 

(  15  minutes. 

Filtered  after 

(26) 

0.1990 

1.0 

.  . 

50 

0.1987 

0.0003- 

cooling;  stood 

15  minutes. 

E. 

(  Filtered  after 

(27) 

0.1990 

2.0 

5.0 

55 

0.1943 

0.0047- 

<  cooling  ;  stood 

(  15  minutes. 

(28) 
(29) 

0.1990 
0.1990 

2.0 
2.0 

5.0 
6.0 

55 
55 

0.1969 
0.1973 

0.0021- 
0.0017- 

Stood  2£  hours. 
Stood  6  hours. 

(30) 

0.1990 

2.0 

5.0 

55 

0.1989 

0.0001- 

Stood  16  hours. 

Some  experiments  were  made  to  show  the  time  necessary 
for  the  complete  precipitation,  both  in  the  presence  and  absence 
of  nitric  acid.  Above  is  the  record  of  such  work. 

The  results  in  section  D  would  seem  to  show  that  a  solution 
containing  copper  may  be  precipitated  hot  as  the  oxalate  and 
filtered  either  hot  or  after  cooling  with  a  very  slight  loss. 
VOL.  ii.  — 23 


354 


VOLUMETRIC  ESTIMATION  OF  COPPER. 


Tests  of  the  filtrates  made  with  potassium  ferrocyanide  con- 
firmed these  results.  When  nitric  acid  is  present,  however, 
the  mixture  must  stand  after  the  addition  of  the  precipitant. 
In  section  E  the  gradual  decrease  of  the  minus  error  is  noticed, 
as  the  time  of  standing  is  extended,  the  precipitation  being 
practically  complete  upon  standing  over  night. 

Separation  from  Cadmium. 

Bornemann  *  has  used  nitric  acid  for  a  rough  separation  of 
copper  from  cadmium.  This  method  was  tried  for  a  quantita- 
tive separation  in  the  presence  of  6-10  per  cent  strong  nitric 
acid.  The  results  are  found  in  section  F  of  the  table  to  follow. 
Experiments  (33)-(35)  stood  six  hours  before  filtering.  Ex- 
periments (36)  and  (37)  stood  over  night.  Copper  is  separated 
from  more  than  twice  its  weight  of  cadmium,  and  the  results 
are  accurate. 

TABLE   III  (continued). 


Exp. 

CuO 

taken  as 
CuS04. 

Element 
from  which 
copper  was 
separated. 

Oxalic 
acid. 

HN03 

(sp.  gr. 
1.40). 

Volume 
at  precip- 
itation. 

CuO 
found. 

Error. 

CdO  taken 
as  CdS04. 

F. 

(33) 
(34) 
(35) 
(36) 
(37) 

gTDl. 

0.1990 
0.1990 
0.1990 
0.1990 
0.1990 

grm. 
0.10 

0.20 
0.30 
0.40 
0.50 

griii. 

2.0 
2.0 
2.0 
2.0 
2.0 

cm» 
5.0 
5.0 
5.0 
5.0 
6.0 

cm8 
60 
65 
70 
76 
80 

grm. 
0.1983 
0.1987 
0.1987 
0.1994 
0.1996 

gnu. 

0.0007- 
0.0003- 
0.0003- 
0.0004+ 
0.00064- 

As2O3  taken 
as  NagAsOs. 

G. 

(38) 
(39) 
(40) 
(41) 

(42) 
(43) 

0.1990 
0.1990 
0.1990 
0.1990 
0.1990 
0.1990 

0.10 
0.20 
0.50 
0.10 
0.20 
0.60 

2.0 
2.0 
2.0 
2.0 
2.0 
2.0 

5.0 
5.0 
5.0 

65 
60 
75 

60 

75 

85 

0.1991 
0.1987 
0.1986 
0.1994 
0.1992 
0.1995 

0.00014- 
0.0003- 
0.0004- 
0.00044- 
0.00024- 
0.00054- 

As,OB  taken 
as  H2KAsO4. 

H. 

(44) 
(45) 
(46) 
(47) 
(48) 
(49) 

0.1990 
0.1990 
0.1990 
0.1990 
0.1990 
0.2030 

0.10 
0.20 
0.10 
0.20 
0.30 
0.30 

2.0 
2.0 
2.0 
2.0 
2.0 
3.0 

5.0 
5.0 
5.0 

5.0 

60 
70 
65 
75 

85 
85 

0.1985 
0.1990 
0.1990 
0.1992 
0.1985 
0.202f> 

0.0005- 
0.0000 
0.0000 
0.00024- 
0.0005— 
0.0004- 

*  Loc.  cit. 


VOLUMETRIC  ESTIMATION  OF  COPPER. 


355 


TABLE  m  (continued). 


Exp. 

Cu 
taken  as 
CuSO4. 

Element 
from  which 
copper  was 
separated. 

Oxalic 
acid. 

HN03 

'So8)'' 

Volume 
at  precip- 
itation. 

Cu 
found. 

Error. 

Sn  taken  as 
SnClj  +  HCl. 

I. 

grm. 
(50) 
(51) 
(51  a) 
(52) 
(53) 
(54) 
(55) 

grin. 
0.1590 
0.1590 
0.1590 
0.1590 
0.1590 
0.1590 
0.1590 

grin. 

0.0468 
0.0936 
0.0936 
0.0936 
0.1873 
0.2809 
0.2809 

grm. 
2.0 
2.0 
2.0 
2.0 
2.0 
2.0 
3.0 

cm» 
5.0 
6.0 
5.0 
5.0 
5.0 
5.0 
5.0 

cm8 
65 
60 
60 
60 
65 
70 
75 

grm. 
0.1681 
0.1603 
0.1591 
0.1694 
0.1603 
0.1914 
0.1988 

grm. 
0.0009- 
0.0013+ 
0.0001+ 
0.0004+ 
0.0013+ 
0.0324+ 
0.0398+ 

Sn  taken 

as  SnCl4. 

K. 

(56) 
(57) 
(58) 
(59) 

0.1590 
0.1590 
0.1590 
0.1590 

0.10 

0.10 
0.20 
0.50 

2.0 
2.0 
2.0 
2.0 

5.0 

5.6 
6.0 

65 
55 
55 
60 

0.1581 
0.1565 
0.1577 
0.1562 

0.0009- 
0.0026- 
0.0013- 
0.0028- 

CuO  taken 
as  CuSO4. 

Fe«O3  taken 
as  Fe(N03)3. 

L. 

CuO  found. 

(60) 
(61) 
(62) 
(63) 
(64) 
(65) 
(66) 

0.1990 
0.1990 
0.1990 
0.1990 
0.1990 
0.1990 
0.1990 

0.136 
0.272 
0.364 
0.544 
0.272 
0.544 
0.218 

2.0 
2.0 
2.0 
2.0 
2.0 
2.0 
2.0 

6.0 
6.0 
5.0 
5.0 

2.6 
2.0 

60 
60 
60 
65 
60 
60 
65 

0.1987 
0.1983 
0.1988 
0.1971 
0.1995 
0.1998 
0.1999 

0.0003- 
0.0007- 
0.0002- 
0.0019- 
0.0005+ 
0.0008+ 
0.0009+ 

ZnO  taken 
as  ZnSO4. 

M. 

(67) 
(68) 
(69) 
(70) 

0.1990 
0.1990 
0.1990 
0.1990 

0.028 
0.057 
0.057 
0.085 

2.0 
2.0 
2.0 
2.0 

5.0 
6.0 
6.0 
6.0 

60 
65 
65 
70 

0.2007 
0.2008 
0.2008 
0.2036 

0.0017+ 
0.0018+ 
0.0018+ 
0.0045+ 

Separation  from  Arsenic,  in  Both  Conditions  of  Oxidation. 

For  the  separation  of  arsenic,  arsenious  oxide  dissolved  in 
sodium  carbonate,  and  di-hydrogen  sodium  arseniate  were  the 
forms  of  arsenic  used.  The  results  are  accurate  and  are  given 
in  sections  G  and  H  of  the  table.  In  experiments  (38)-(40) 
and  (44)-(45)  no  nitric  acid  was  added.  While  the  presence 
of  the  nitric  acid  is  not  necessary  for  the  separation  of  the 
copper  from  the  arsenic,  still  the  filtration  in  the  absence  of 


356  VOLUMETRIC  ESTIMATION  OF  COPPER. 

the  nitric  acid  is  so  slow  as  to  be  objectionable.  The  presence 
of  the  nitric  acid  causes  the  precipitate  to  come  down  in  a 
coarser  condition,  and  in  such  condition  it  filters  easily  and  is 
capable  of  being  washed  quickly. 

Separation  from  Tin,  in  Both  Conditions  of  Oxidation. 

For  the  separation  of  copper  from  tin  a  preparation  of 
stannous  chloride  (20  cm3  giving  0.3746  grm.  metallic  tin  by 
the  battery)  containing  sufficient  hydrochloric  acid  to  prevent 
deposition  of  oxy-salts  was  used.  The  solution  of  stannic 
chloride  contained  1.0  grm.  metallic  tin  to  every  10  cm3,  and 
was  used  without  hydrochloric  acid.  The  results  of  the  work 
are  found  in  sections  I  and  K  of  the  table.  The  experiments 
go  to  show  that  while  copper  may  be  separated  from  small 
amounts  of  tin  as  stannous  chloride  yet  there  is  a  limit  to  the 
amount  of  tin  which  may  be  present.  One-tenth  of  a  gram  of 
metallic  tin  is  the  largest  amount  that  can  be  present,  with 
0.15  gm.  copper  oxide  taken  as  the  sulphate,  without  significant 
error.  Practically  the  same  statement  can  be  made  of  the 
separation  of  copper  from  tin  taken  as  stannic  chloride. 
Experiment  (57)  shows  a  greater  loss  of  copper  when  the 
nitric  acid  is  omitted. 

Separation  of  Copper  from  Iron. 

A  solution  of  ferric  nitrate  was  used  for  the  work  on  the 
separation  of  copper  from  iron.  Low  results  were  obtained 
when  a  solution  of  ferrous  or  ferric  sulphate  was  used  as  the 
source  of  iron.  The  results  of  the  experiments  are  recorded 
in  section  L  of  the  table,  and  show  that  0.20  grm.  copper  oxide 
as  the  sulphate  may  be  separated  from  0.2-0.3  grm.  iron  oxide 
taken  as  the  nitrate.  In  experiment  (64)  a  good  result  was 
obtained  when  no  nitric  acid  was  present,  save  that  added  in 
combination  with  the  iron.  A  comparison  of  experiments 
(63)  and  (65)  shows  that  it  is  best  to  avoid  the  use  of  large 
amounts  of  nitric  acid  when  the  larger  amounts  of  ferric 
nitrate  are  present. 


VOLUMETRIC  ESTIMATION  OF  COPPER.          357 

For  a  practical  application  of  the  above  separation  of  copper 
from  iron  a  convenient  amount  of  finely  ground  chalcopyrite 
(0.5  gnn.)  was  roasted  2-3  hours  in  a  porcelain  crucible  until 
all  sulphur  was  driven  off,  washed  into  a  beaker,  strong  nitric 
acid,  about  5  cm3,  was  added  and,  with  the  beaker  covered, 
allowed  to  evaporate  slowly  on  a  hot  plate,  nearly  to  dryness. 
A  little  dilute  nitric  acid  was  added,  the  solution  was  filtered, 
the  residue  was  washed  with  water  containing  dilute  nitric 
acid,  the  filtrate,  about  50  cm3  in  volume,  was  precipitated 
with  2.0  grm.  oxalic  acid,  and  the  precipitate  was  estimated 
after  standing  12-16  hours,  as  previously  described.  The 
washing  with  water  acidified  with  nitric  acid  is  important, 
because  the  finely  ground  ferric  oxide  remaining  undissolved 
passes  through  the  filter  when  washed  with  water  alone,  but 
gives  no  trouble  if  the  water  be  acidic.  The  results  of  two 
estimations  are  here  given. 


Chalcopyrite. 

Copper  found 
by  battery. 

Copper  found  by 
oxalate  method. 

Difference. 

grm. 

% 

% 

% 

0.5000 

31.00 

30.92 

0.08- 

1.0000 

31.00 

31.25 

0.25+ 

Separation  of  Copper  from  Zinc. 

The  separation  of  copper  from  zinc  was  not  altogether 
successful  owing  to  the  tendency  of  the  zinc  oxalate  to  come 
down  with  the  copper  oxalate.  Some  experiments  are  given 
in  section  M  of  the  table. 

The  separations  of  copper  from  bismuth  and  antimony  were 
unsuccessful. 

The  work  may  be  briefly  summarized  as  follows :  Copper 
exceeding  in  amount  the  equivalent  of  0.0128  grm.  of  the 
oxide  to  50  cm3  of  solution  as  the  sulphate  may  be  separated 
completely,  even  in  the  presence  of  a  moderate  amount  of 
strong  nitric  acid,  by  the  addition  of  sufficient  amount  of 
oxalic  acid. 

Copper  may  be  separated  from  cadmium,  arsenic,  iron,  and 
small  amounts  of  tin,  when  precipitated  by  oxalic  acid  in 


358  VOLUMETRIC  ESTIMATION  OF  COPPER. 

a  volume  of  50  cm3  containing  5  cm3  strong  nitric  acid. 
Inasmuch  as  the  completeness  of  precipitation  of  the  copper 
depends  upon  the  presence  of  a  certain  minimum  amount 
of  the  copper  salt  this  method  is  not  applicable  when  the 
amount  of  copper  falls  below  0.0128  grm.  of  the  oxide  to  50 
cm3  of  solution. 


XLVI 

THE    SULPHOCYANIDES    OF    COPPER    AND    SIL- 
VER IN   GRAVIMETRIC  ANALYSIS. 

BY  K.  G.  VAN   NAME  * 

Cuprous  Sulphocyanide. 

As  early  as  1854  attention  was  drawn  by  Rivotf  to  the 
possibility  of  estimating  copper  gravimetrically  by  weighing 
as  cuprous  sulphocyanide,  and  to  the  advantages  which  the 
process  afforded  in  separating  copper  from  other  metals. 
Rivot's  procedure  consisted  in  dissolving  the  substance  to 
be  analyzed  in  hydrochloric  acid,  reducing  copper  with  hypo- 
phosphorous  or  sulphurous  acid,  and  precipitating  with  potas- 
sium sulphocyanide.  The  precipitate  dried  at  a  moderate 
temperature  was  weighed  as  cuprous  sulphocyanide  and  then 
as  a  control  converted  by  ignition  with  sulphur  into  cuprous 
sulphide  and  weighed  in  that  condition. 

In  his  well  known  work  upon  quantitative  analysis  Fre- 
senius  in  one  place  J  denies  the  practicability  of  the  direct 
weighing  of  copper  as  cuprous  sulphocyanide  on  account  of 
the  tendency  of  the  latter  to  hold  water  even  when  heated  to 
the  temperature  of  incipient  decomposition.  As  authority 
for  this  statement  he  cites  Claus,§  who  found  3  per  cent  of 
water  in  the  precipitate  after  drying  at  115°,  and  Meitzendorff, 
who  gave  the  percentage  of  water  under  the  same  conditions 
as  1.54. 

On  a  later  page  of  the  same  volume,  ||  however,  Fresenius, 
after  a  trial  of  the  process  which  gave  99.66  per  cent  of  the 

*  From  Am.  Jour.  Sci.,  x,  451.  t  Compt.  rend.,  xxxviii,  868. 

J  Quant.  Anal.,  6.  Aufl.,  i,  187.  §  L.  Gmelin,  Handbuch,  iv,  472. 

||  Quant.  Anal,  6.  Aufl.,  i,  335. 


360      SULPHOCYANIDES  OF  COPPER  AND  SILVER 

theory  for  copper,  concludes  that  the  method  is  practicable 
although  apt  to  give  low  results,  particularly  in  the  presence 
of  free  acid. 

The  process  was  again  recommended  in  1878  by  Busse,* 
who  had  employed  it  for  the  estimation  of  copper,  both 
alone  and  in  the  presence  of  iron,  nickel,  zinc,  and  arsenic, 
obtaining  results  very  near  the  theory  and  plainly  comparable 
with  the  figures  obtained  by  afterwards  igniting  the  cuprous 
sulphocyanide  with  sulphur  in  hydrogen. 

In  spite  of  the  evident  advantages  for  certain  purposes  of 
Rivot's  method  over  other  modes  of  determining  copper,  it 
has  never  come  into  general  use.  The  chief  reason  for  this 
has  apparently  been  the  difficulty  and  inaccuracy  attendant 
upon  the  weighing  of  the  precipitate  upon  dried  paper 
filters,  a  process  which  can  hardly  be  depended  upon  unless 
managed  with  extreme  care. 

In  the  experiments  to  be  described  this  difficulty  was 
avoided  by  performing  the  filtering  and  weighing  upon  asbes- 
tos in  a  perforated  platinum  crucible.  The  method  of  con- 
ducting a  determination  was  as  follows :  A  suitable  quantity 
of  a  standard  copper  sulphate  solution  was  run  from  a  burette, 
diluted  to  a  convenient  volume,  a  few  cubic  centimeters  of  a 
saturated  solution  of  ammonium  bisulphite  added,  and  the 
copper  precipitated  by  an  excess  of  ammonium  sulphocyanide. 
The  precipitate  was  allowed  to  settle,  collected  upon  asbestos 
in  a  weighed  crucible,  washed  with  cold  water,  and  dried  at 
110°  until  no  further  loss  of  weight  took  place. 

In  Table  I  are  given  the  results  of  a  number  of  determi- 
nations made  in  this  way.  The  copper  sulphate  solution 
was  made  up  exactly  decinormal  and  the  standard  confirmed 
electrolytically.  As  the  ammonium  sulphocyanide  solution 
was  slightly  above  decinormal,  13  cm3  represent  a  small  excess 
(about  one'  cubic  centimeter)  above  the  amount  theoretically 
required  to  precipitate  25  cm3  of  the  copper  sulphate  solution. 
The  ammonium  bisulphite,  which  had  been  recently  prepared 
by  saturating  aqueous  ammonia  with  sulphur  dioxide,  was 

*  Zeitschr.  anal.  Chem.,  xvii,  63. 


IN  GRAVIMETRIC  ANALYSIS. 


361 


always  used  in  sufficient  quantity  to  give  the  liquid  a  strong 
and  permanent  odor  of  the  latter. 

TABLE  I. 

25  cm8  of  ^  CuSO4  solution,  equivalent  to  0.0795  grm.  Cu,  taken  for  each 

experiment. 


Exp. 

H2S04. 
concentrated. 

HNH4SOS 
sat.  sol. 

NH4SCN 
approz. 

ff 

Final 
volume. 

Time  of 
standing. 

Cu 
found. 

Error. 

TIT 

cms 

cm3 

cm3 

cm3 

hrs. 

grui. 

grm. 

(1) 

None. 

5 

13 

68 

| 

0.0795 

0.0000 

(2) 

None. 

3 

13 

66 

48 

0.0793 

0.0002- 

(3) 

None. 

3 

25 

78 

i 

0.0796 

0.00014 

(4) 

None. 

3 

25 

78 

12 

0.0796 

0.0001+ 

(5) 

1.5 

10 

13 

85 

12 

0.0792 

0.0003- 

(6) 

1.6 

8 

13 

105 

48 

0.0785 

0.0010- 

(7) 

1.5 

3 

25 

85 

4 

0.0783 

0.0012- 

(8) 

1.5 

5 

25 

85 

21 

0.0795 

0.0000 

(9) 

5 

5 

25 

85 

3 

0.0797 

0.0002+ 

(10) 

15 

10 

25 

115 

21 

0.0793 

0.0002- 

HCl 

concentrated. 

(11) 

10 

6 

25 

100 

20 

0.0795 

0.0000 

(12) 

25 

10 

25 

100 

28 

0.0784 

0.0011- 

When  there  is  no  free  acid  present  the  time  of  standing 
before  nitration  and  the  amount  of  the  excess  of  ammonium 
sulphocyanide  are  practically  without  effect,  as  experiments 
(1)  to  (4)  of  the  table  show. 

Experiments  (5)  to  (10)  were  carried  out  in  the  presence  of 
various  amounts  of  free  sulphuric  acid  up  to  12  per  cent  of 
the  total  volume  of  liquid.  The  acid,  at  least  within  this 
limit,  does  not  exert  a  sufficient  solvent  effect  upon  the  cuprous 
sulphocyanide  to  interfere  materially  with  the  accuracy  of  the 
process,  but  it  retards  the  precipitation,  making  it  necessary  to 
increase  the  time  of  standing  before  filtering  in  proportion  to 
the  amount  of  acid  present.  In  several  of  these  determina- 
tions the  precipitation  was  visibly  incomplete  even  after  sev- 
eral hours'  standing.  This  effect  of  the  acid,  however,  hardly 
shows  in  the  results  of  the  table  because  the  standing  was 
always  prolonged  until  the  copper  appeared  to  be  all  down 
before  filtering. 


362       SULPHOCYANIDES   OF  COPPER  AND  SILVER 


The  low  result  of  Experiment  (7)  was  probably  due  chiefly 
to  incomplete  precipitation,  although  (9)  shows  that  even  with 
a  much  larger  amount  of  acid  precipitation  may  be  complete 
within  three  hours.  In  general,  however,  it  is  safer  to  allow 
ample  time  (twelve  hours  or  more)  for  the  precipitation  when 
there  is  much  free  acid  present. 

Comparison  of  Experiments  (5)  and  (6),  for  which  only  a 
bare  excess  of  ammonium  sulphocyanide  was  used,  with  (7) 
to  (12)  shows  an  apparent  advantage  in  the  larger  excess  in  the 
presence  of  acid.  Hydrochloric  acid,  judging  from  the  results 
of  (11)  and  (12),  has  no  greater  disturbing  influence  than  sul- 
phuric acid,  although  in  (12),  where  the  concentrated  acid 
constituted  one-fourth  of  the  entire  volume,  there  was  appar- 
ently a  slight  solvent  action.  The  filtrate  from  this  determi- 
nation when  concentrated  to  about  25  cm8  and  treated  with 
potassium  ferrocyanide  gave  a  strong  test  for  copper,  as  did 
also  the  filtrate  from  (6).  Several  of  the  other  nitrates  were 
tested  in  the  same  way,  but  none  showed  more  than  an  insig- 
nificant trace  of  copper.  The  nitrate  of  (7),  however,  was 
not  tested. 

Table  II  contains  the  results  of  a  series  of  experiments  con- 
ducted as  before,  except  that  larger  amounts  of  copper  were 
employed.  The  copper  sulphate  solution  was  approximately 
f  and  standardized  by  the  battery.  The  solution  of  ammo- 
nium sulphocyanide  was  the  same  previously  used  and  a  con- 
siderable excess  was  employed  in  every  determination.  More 

TABLE  II. 


Bxp. 

Cu 
taken. 

H2S04 
concen- 
trated. 

NH4SCN 
approx. 
ff 
TS' 

Final 
volume. 

Cu2S2(CN)2 
found,  calcu- 
lated as  Cu. 

Error. 

Cuin 
filtrate. 

| 

(3) 
(4) 

grm. 

0.3175 
0.3176 
0.3175 
0.3175 

cms 

None. 
None. 
None. 
10 

cm8 
60 
60 
60 
100 

cm8 
500 
500 
500 
500 

grm. 
0.3176 
0.3177 
0.3176 
0.3176 

grm. 
0.0001+ 
0.0002+ 
0.0001+ 
0.0000 

None. 
None. 
None. 
None. 

(5) 

0.3175 

HC1 
cone. 

100 

600 

0.3165 

0.0010- 

Distinct. 

20 

IN,  GRA  VIMETRIC  ANAL YSIS.  363 

than  twice  the  amount  theoretically  required  was  used  in  every 
case  where  free  acid  was  present,  and  at  least  twenty  hours 
were  allowed  for  the  precipitation,  which  was  made  in  cold, 
and  as  the  table  shows,  rather  dilute  solutions.  If  the  solu- 
tion is  too  concentrated  the  copper  is  apt  to  be  thrown  down 
in  a  finely  divided  condition,  making  it  hard  to  filter. 

The  time  required  to  dry  the  cuprous  sulphocyanide  at  110° 
is  in  general  from  two  to  three  hours.  Heating  much  longer 
than  this  is  not  to  be  recommended,  as  a  gradual  increase  in 
weight  begins  to  take  place,  as  is  shown  by  the  following 
example,  which  gives  a  series  of  weights  of  the  same  precipi- 
tate at  different  stages. 

Cu,S8(CN),.          Calculated  as  Cu. 
gnu.  gnu. 

After  2  hours  at  110°  .     .     .  0.6060  0.3167 

«     4        "         «  ...  0.6059  0.3167 

«    19        "         «  .    .     .  0.6067  0.3171 

«   23        «         "  ...  0.6069  0.3172 

This  tendency  to  increase  in  weight  is,  however,  usually  less 
marked  than  in  the  above  example,  and  in  any  case  need  not 
interfere  materially  with  the  accuracy  of  the  process  unless  the 
drying  is  prolonged  far  beyond  the  necessary  length  of  time. 

The  method  is  easily  handled  and,  as  the  results  of  Tables  I 
and  II  show,  is  capable  of  considerable  accuracy.  From  the 
nature  of  the  process  it  is  evident  that  it  is  much  less  likely  to 
be  interfered  with  by  the  presence  of  other  metals  than  the 
other  gravimetric  methods  for  copper,  and  may  therefore  be 
directly  applied  with  good  results  in  many  cases  where  the  use 
of  the  electrolytic  or  the  oxide  method  would  involve  a  pre- 
vious separation. 

Silver  Sulphocyanide. 

The  sulphocyanide  of  silver,  unlike  that  of  copper,  is  solu- 
ble in  an  excess  of  ammonium  or  alkali  sulphocyanides  and 
this  fact  prevents  the  use  of  the  latter  to  precipitate  silver  for 
gravimetric  estimation.  The  reverse  process,  however,  the 
precipitation  of  a  soluble  sulphocyanide  by  an  excess  of  silver 


364       SULPHOCYANIDES  OF  COPPER  AND  SILVER 

nitrate,  as  will  be  shown  by  the  experiments  to  be  described, 
furnishes  a  convenient  means  of  standardizing  sulphocyanide 
solutions  and  in  general  for  estimating  sulphocyanic  acid. 

When  freshly  precipitated  the  sulphocyanide  of  silver  resem- 
bles the  chloride  in  appearance,  but  when  allowed  to  stand  a 
few  hours  becomes  finely  granular  and  is  very  easily  filtered 
and  washed.  It  may  be  safely  dried  to  a  constant  weight  upon 
an  asbestos  filter  at  110° -120°,  but  at  a  somewhat  higher  tem- 
perature is  decomposed,  leaving  a  residue  of  silver  sulphide. 

The  determinations  which  are  tabulated  below  were  made  as 
follows.  Portions  of  25  cm3  of  an  approximately  decinormal 
solution  of  ammonium  sulphocyanide  were  measured  from  a 
burette,  diluted  with  100  cm3  of  water  and  silver  nitrate  added 
in  excess.  The  precipitate  was  collected  upon  asbestos  in  a 
platinum  crucible,  washed  with  cold  water  and  dried  to  a  con- 
stant weight  at  115°  the  drying  requiring  usually  between  two 
and  three  hours. 

The  filtering  is  facilitated  by  allowing  a  few  hours  for  the 
precipitate  to  settle  ;  but  this  is  by  no  means  essential,  as  it  is 
easy  with  a  little  care  to  obtain  a  clear  filtrate  even  when  the 
filtering  is  performed  at  once. 

The  solution  of  ammonium  sulphocyanide  was  prepared 
from  a  pure  salt,  especially  tested  and  found  free  from  choride. 
This  point  is  of  importance,  as  chlorine  is  a  common  impurity 
and  its  presence  in  any  considerable  quantity  will  vitiate  the 
results. 

In  order  that  the  effect  of  varying  the  excess  of  silver 
might  be  investigated,  an  approximately  decinormal  solution 
of  silver  nitrate  was  titrated  against  the  ammonium  sulpho- 
cyanide and  the  ratio  between  the  two  solutions  determined. 
This  silver  nitrate  solution  was  used  for  the  first  five  determi- 
nations of  Table  III.  For  the  rest  the  quantity  of  silver  nitrate 
was  not  measured  but  regulated  by  the  eye  alone,  thus  making 
the  conditions  the  same  as  would  be  the  case  in  practical  use 
of  the  method. 

These  results  are  as  uniform  as  could  be  expected,  considering 
the  variations  which  would  be  produced  by  even  very  small 


GRAVIMETRIC  ANALYSIS. 


365 


TABLE  in. 

Final  volume  of  liquid  150  cm8. 
25  cm8  of  NH4SCN  sol.  equivalent  to  25.15  cm8  of  AgN08  sol. 


Exp. 

NH3SCN. 

AgN08. 

Excess  of 
AgNOs. 

AgSCN 
found. 

cm3 

cm3 

cm3 

grm. 

(1) 

25 

25.3 

0.15 

0.4372 

(2) 

25 

25.3 

0.15 

.0.4376 

8 

25 
25 

25.4 
25.4 

0.25 
0.25 

0.4373 
0.4375 

(5) 

25 

30.4 

5.25 

0.4382 

(6) 

25 

Kough  excess. 

0.4366 

(7) 

25 

Rough  excess. 

0.4381 

(8) 

25 

Rough  excess. 

0.4873 

(9) 

25 

Rough  excess. 

0.4372 

(10 

25 

Rough  excess. 

0.4369 

errors  in  measuring  out  25  cm3  of  decinormal  sulphocyanide 
solution.  It  is  moreover  clearly  shown  that  there  is  no 
difference  in  the  results  whether  a  bare  excess  or  a  moderately 
large  excess  of  the  silver  nitrate  is  used. 

The  mean  of  the  values  in  the  last  column  is  0.4374,  which  is 
equivalent  to  0.2006  grm.  of  ammonium  sulphocyanide  for 
every  25  cm3  of  the  solution. 

The  standard  of  the  sulphocyanide  solution  was  also 
determined  volumetrically  by  Volhard's  process.  The  mean  of 
four  titrations  carried  out  with  great  care  against  a  standard 
silver  nitrate  solution  gave  as  the  standard  0.2003  grm.  of 
ammonium  sulphocyanide  for  25  cm3  of  solution.  This  differ- 
ence between  the  standards  as  determined  by  the  two  methods 
(one  part  in  670)  is  much  less  than  the  variations  which 
frequently  appear  between  successive  determinations  by 
Volhard's  method,  under  like  conditions  as  to  strength  of 
solutions  and  amounts  used.  It  is  about  equal  to  the  error 
that  would  be  produced  in  a  single  volumetric  determination 
by  a  mistake  of  one  drop  in  measuring  one  of  the  solutions,  or 
of  one-half  drop  in  the  same  direction  on  each. 

It  is  therefore  evident  that  the  standard  of  a  sulphocyanide 
solution  obtained  in  the  above  way  may  be  applied  directly  to 


366       SULPHOCYANIDES  OF  COPPER  AND  SILVER 

the  estimation  of  unknown  amounts  of  silver  by  Volhard's 
method  without  sensible  error. 

To  remove  a  possible  doubt  as  to  whether  the  silver 
sulphocyanide  dried  at  115°  was  entirely  free  from  water,  a 
number  of  electrolytic  determinations  of  the  silver  contained 
in  the  previously  weighed  precipitates  of  Table  III  were  made 
in  the  following  way. 

The  perforated  platinum  crucible  containing  the  silver 
sulphocyanide  and  asbestos  was  hung  in  a  loop  of  heavy 
platinum  wire  and  served  as  the  anode.  For  the  cathode  a 
deep  platinum  dish  of  about  200  cm3  capacity  was  used.  An 
ammoniacal  solution  of  potassium  cyanide  was  employed  as  the 
electrolyte  and  gave  the  best  results  when  made  up  by  dissolving 
2  grm.  of  potassium  cyanide  in  15  cm3  of  strong  ammonia  and 
15  cm3  of  water.  The  crucible  which  served  as  the  anode  was 
filled  with  this  solution  in  full  strength,  and  the  remainder 
was  put  into  the  platinum  dish  and  diluted  to  the  required 
volume  with  water.  In  this  medium  the  silver  sulphocyanide 
is  slowly  dissolved  and  diffuses  through  the  asbestos  felt  into 
the  space  between  the  electrodes  where  the  silver  is  deposited 
in  the  usual  way.  This  diffusion,  is,  however,  aided  but  little 
if  at  all  by  the  current,  and  there  is  a  tendency  for  traces  of 
the  silver  to  remain  behind  in  the  crucible.  The  current  density 
employed  was  about  0.0012  ampere  per  square  centimeter 
of  cathode  surface  and  the  time  about  twelve  hours.  After 
weighing  the  silver  deposited,  it  was  dissolved  in  nitric  acid, 
precipitated  by  hydrochloric  acid  and  weighed  again  as  the 
chloride,  giving  a  check  upon  the  results. 

Seven  of  the  ten  determinations  of  Table  III  were  thus 
treated,  but  owing  to  the  imperfections  of  the  process  the 
results  were  all  slightly  low,  the  worst  showing  a  deficiency 
of  0.0025  grm.  of  silver,  an  error  of  less  than  0.9  per  cent.  The 
results  of  the  two  best  of  these  determinations  given  below  are, 
however,  sufficient  to  prove  the  point  hi  question,  namely 
that  the  silver  sulphocyanide  dried  at  115°  has  the  theoretical 
constitution  and  contains  no  water.  The  numbers  are  those 
under  which  the  determinations  appear  in  Table  III. 


/.AT  GRAVIMETRIC  ANALYSIS. 


367 


Bxp. 

AgSCN 
taken. 

Calculated 
as  Ag. 

Ag  found 
by  battery. 

Error. 

Weighed 
AgCl. 

Calculated 
asAg. 

Error. 

(4) 
(10) 

grm. 

0.4375 
0.4369 

gnu. 
0.2844 
0.2840 

grm. 
0.2839 
0.2838 

grm. 
0.0005- 
0.0002- 

grm. 
0.3765 
0.3761 

grm. 
0.2834 
0.2831 

grm. 
0.0010- 
0.0009- 

It  is  clear  therefore  that  the  estimation  of  sulphocyanides  by 
precipitation  with  silver  nitrate  and  direct  weighing  of  the 
precipitate  is  wholly  permissible.  The  method  is  extremely 
simple  and,  as  has  been  shown,  the  results  are  quite  accurate. 


XLVII 

ON  THE  ESTIMATION  OF  CAESIUM  AND  EUBIDIUM 
AS  THE  ACID  SULPHATES,  AND  OF  POTASSIUM 
AND  SODIUM  AS  THE  PYROSULPHATES. 

BY  PHILIP  E.  BROWNING. 

BUNSEN  *  is  authority  for  the  statement  that  the  acid  sulphate 
of  rubidium  does  not  lose  sulphuric  acid  at  a  heat  approaching 
redness.  It  is  stated  in  the  literature  f  that  the  acid  sulphates 
of  caesium  and  rubidium  when  subjected  to  a  low  red  heat 
pass  into  the  form  of  the  pyrosulphates. 

R.  Weber  J  found  that  by  treating  the  dry  sulphates  of 
potassium,  caesium,  rubidium,  and  thallium  with  sulphuric 
anhydride  in  a  closed  tube  and  heating  on  a  water  bath  two 
layers  separated.  In  the  lower  layer  he  obtained  crystalline 
bodies  which  proved  to  have  the  constitution  R2O  .  8SO3.  On 
strong  heating  he  obtained  from  these  substances,  bodies  of 
the  form  R20 .  2SO3  and  finally  R2O  .  SO3.  He  also  notes  that 
in  the  case  of  the  caesium  salt  the  removal  of  the  excess  of 
the  sulphuric  anhydride  was  attended  with  greater  difficulty. 

Baum  §  states  that  the  pyrosulphates  of  the  alkalies  may  be 
obtained  by  heating  the  acid  sulphates  under  atmospneric 
pressure  at  low  redness,  or  under  diminished  pressure  at  a 
temperature  between  260°  C.  and  320°  C. 

In  a  recent  paper  ||  from  this  laboratory  I  have  shown  that 
thallium  may  be  estimated  as  the  acid  sulphate  by  evaporating 
a  thallous  salt  in  solution  with  an  excess  of  sulphuric  acid  and 
bringing  the  residue  to  a  constant  weight  at  a  temperature  of 
about  250°  C. 

*  Ann.  Chem.  (Liebig),  cxix,  110. 

t  Graham-Otto,  Lehrbuch  d.  Chem.,  iii,  269,  278. 

|  Ber.  Dtsch.  Chem.  Ges.,  xvii,  2497.      §  Ber.  Dtsch.  Chem.  Ges.  xx,  752. 

||  Am.  Jour.  Sci.,  ix  (1900),  137.    This  volume,  p.  317. 


ESTIMATION  OF  CAESIUM  AND  RUBIDIUM.        369 


The  similarity  which  thallium  bears  in  some  of  its  combi- 
nations to  the  alkaline  metals  suggested  the  study  of  the 
sulphates  of  these  elements  under  the  same  general  conditions 
of  procedure. 

My  first  experiments  were  made  with  a  pure  caesium  salt 
as  follows :  A  weighed  amount  of  the  nitrate  was  placed  in 
a  previously  weighed  platinum  crucible  and  treated  with  an 
excess  of  sulphuric  acid.  The  crucible  was  then  placed  upon 
a  steam  bath  until  the  water  and  nitric  acid  were  largely 
expelled  and  then  removed  to  a  radiator,  consisting  of  a 
porcelain  crucible  fitted  with  a  pipe-stem  triangle  so  arranged 
that  the  bottom  of  the  platinum  crucible  would  be  about 
midway  between  the  top  and  bottom  of  the  porcelain  crucible. 
This  improvised  radiator  was  set  in  an  iron  ring  and  a 
thermometer  so  placed  that  the  mercury  bulb  would  be  on 
a  level  with  the  bottom  and  close  to  the  side  of  the  platinum 
crucible.  An  ordinary  Bunsen  burner  served  as  the  source 
of  heat  and  the  temperature  was  kept  so  far  as  possible 
between  250°  C.  and  270°  C.  After  the  fuming  attending 
the  removal  of  the  large  excess  of  sulphuric  acid  ceased,  the 
crucible  and  contents  were  removed  to  a  desiccator,  and,  after 
being  allowed  to  cool,  weighed.  This  process  of  heating 
was  continued  for  half-hour  periods  until  the  weights  were 
constant.  The  results  shown  in  Table  I  were  obtained  by 
this  method  of  treatment.  In  experiments  (1),  (4),  and  (9) 

TABLE  I. 


Ezp. 

CsN08 
taken. 

CsHS04 
calcu- 
lated. 

First 
constant 
weight. 

Second 
constant 
weight. 

Error 
CsHS04. 

Cs2SO4 
calcu- 
lated. 

Cs2S04 
found. 

Error 
Cs2S04. 

grm. 

grm. 

grm. 

grm. 

grin. 

grm* 

grm. 

grm. 

(1) 

0.1706 

0.2013 

0.2054 

0.2020 

0.0007+ 

(2) 

0.1706 

0.2013 

0.2010 

0.0003- 

t 

(3) 

0.1032 

0.1217 

0.1201 

0.0016- 

(4) 

0.1032 

0.1217 

0.1252 

0.1222 

0.0005+ 

0.0961 

0.0948 

0.0013- 

(5) 

0.1218 

0.1437 

0.1458 

( 

0.0021+ 

0.1130 

0.1118 

0.0012- 

(6) 

0.1214 

0.1435 

0.1430 

0.0005- 

( 

(7) 

0.1214 

0.1435 

0.1422 

, 

0.0013- 

( 

(8) 

0.1150 

0.1356 

0.1330 

. 

0.0026- 

(9) 

0.1050 

0.1245 

0.1272 

0.1248 

0.0003+ 

(10) 

0.1056 

0.1245 

0.1252 

0.0007+ 

VOL.  ii.  —  24 


370       ESTIMATION  OF  CESIUM  AND  RUBIDIUM. 


it  will  be  noticed  that  the  weights  were  constant  somewhat 
above  the  condition  of  the  acid  sulphate,  a  fact  which  would 
go  to  show  a  tendency  on  the  part  of  the  caesium  salt  to  hold 
an  excess  of  sulphuric  acid  over  the  amount  necessary  to 
form  the  ordinary  acid  sulphate.  The  results  show  that  by 
regulating  the  heat  at  a  temperature  between  250°  C.  and 
270°  C.  caesium  may  be  brought,  with  a  fair  degree  of  cer- 
tainty to  the  condition  of  the  acid  sulphate.  As  a  check  upon 
the  results,  the  acid  sulphate  was,  in  a  few  cases,  treated  with 
a  little  ammonium  hydroxide,  the  excess  of  this  was  removed 
upon  a  steam  bath  and  the  neutral  sulphate  obtained  by 
ignition  at  a  red  heat  to  a  constant  weight.  These  determi- 
nations agreed  fairly  well  with  the  theory.  The  same  pro- 
cedure was  followed  with  rubidium,  a  pure  rubidium  chloride 
having  been  chosen  as  the  starting-point.  The  results  are 

TABLE  II. 


Exp. 

RbCl 
taken. 

RbHS04 
calculated. 

RbHS04 
found. 

Error. 

Rb2S04 
calculated. 

Rb2S04 

found. 

Error. 

grm. 

grm. 

grm. 

grm. 

grm. 

grm. 

grm. 

(1) 

0.1252 

0.1889 

0.1878 

0.0011- 

.  . 

(2) 

0.1212 

0.1829 

0.1840 

0.0011+ 

0.1460 

0.1460 

0.0000 

(3) 

0.1230 

0.1856 

0.1850 

0.0006- 

(4) 

0.1230 

0.1856 

0.1858 

0.0002+ 

0.1357 

0.1350 

0.0007- 

(5) 

0.1610 

0.2430 

0.2416 

0.0014- 

0.1777 

0.1772 

0.0005- 

(6) 

0.1360 

0.2052 

0.2032 

0.0020- 

0.1501 

0.1490 

0.0011- 

given  in  Table  II.  No  tendency  was  observed  on  the  part 
of  this  element  to  hold  sulphuric  acid  in  excess  of  the  amount 
necessary  for  the  formation  of  the  acid  sulphate.  When  the 
same  method  was  applied  to  sodium  and  potassium  salts, 
pure  chlorides  being  used  as  the  starting-point,  a  somewhat 
different  result  was  obtained,  in  that  the  weight  of  the 
final  product  appeared  to  indicate  the  formation  of  the 
pyrosulphate.  The  results  given  in  Tables  III  and  IV,  in 
which  the  sodium  and  potassium  salts  are  calculated  as 
pyrosulphates,  are  sufficiently  satisfactory  for  purposes  of 
quantitative  estimation.  As  in  the  case  of  the  caesium  and 
rubidium  salts,  a  number  of  determinations  as  the  neutral 


ESTIMATION  OF  CAESIUM  AND  RUBIDIUM.        371 
TABLE  m. 


Exp. 

KC1 
taken. 

K2S,07 
calculated. 

K,8207 
found. 

Error. 

K,S04 
calculated. 

Sffi 

Error. 

gnu. 

grm. 

grm. 

grin. 

grm. 

grm. 

grm. 

(1) 

0.2172 

0,3704 

0.3698 

0.0006- 

2 

0.1706 

0.2909 

0.2886 

0.0023— 

0.1993 

0.1972 

0.0021- 

(3) 

0.1192 

0.2032 

0.2022 

0.0010- 

0.1393 

0.1381 

0.0012- 

(4) 

0.1074 

0.1830 

0.1823 

0.0007- 

(5) 

0.1096 

0.1868 

0.1860 

0.0008- 

•  • 

TABLE  IV. 


Exp. 

NaCl 
taken. 

Na28207 
calculated. 

Na2S207 
found. 

Error. 

Na.jS04 
calculated. 

NajSO* 
found. 

Error. 

(1) 

(2) 
(3) 
(4) 

grin* 

0.1042 
0.1028 
0.1093 
0.1402 

grm. 

0.1978 
0.1952 
0.2075 
0.2662 

grm. 
0.1972 
0.1952 
0.2065 
0.2651 

grm. 

0.0006- 
0.0000 
0.0010- 
0.0011- 

gnu. 

0.1266 

0.1328 
0.1703 

grm. 

0.1254 

0.1320 
0.1696 

grm. 

0.0012- 

0.0008- 
0.0007- 

sulphate  were  made  by  ignition  of  the  sodium  and  potassium 
pyrosulphates,  with  results  which  are  recorded.  In  Table  V, 
two  determinations  are  recorded,  in  one  of  which  the  caesium 
and  rubidium  salts  were  treated  together  and  in  the  other 
the  sodium  and  potassium  salts. 

TABLE  V. 


Exp. 

RbCl  +  CsNOj 
taken. 

RbHS01  + 
CsHS04 
calculated. 

RbHS04  + 
CsHS04 
found. 

Error. 

Rb2S04  + 
Cs2SO4 
calculated. 

Rb2S04  + 
CB*S04 
found. 

Error. 

(1) 

grm. 

(RbCl  0.1428  ) 
{CsN03  0.1264  f 

grm. 

0.3646 

grm. 

0.3666 

grm. 
0.0020+ 

grm. 
0.2749 

grm. 

0.2752 

griii. 
0.0003+ 

NaCl  +  KC1 
taken. 

^80,+ 
MLftfff 
calculated. 

Na2S,07  + 
K2S2O7 
found. 

Error. 

^A 

calculated. 

Na28g4ik 

K2SO4 
found. 

Error. 

(2) 

(NaCl  0.1233   ) 
|KCI  0.1340   ] 

0.4627 

0.4630 

0.0003+ 

0.3062 

0.3040 

0.0022- 

An  application  of  this  general  method  to  a  lithium  salt 
gave  no  evidence  of  the  existence  of  a  stable  acid  sulphate 
or  pyrosulphate. 


372       ESTIMATION  OF  CAESIUM  AND  RUBIDIUM. 

The  results  may  be  summed  up  as  follows :  —  Csesium  and 
rubidium  salts  of  volatile  acids  when  treated  with  sulphuric 
acid  in  excess  and  brought  to  a  constant  weight  at  a  tem- 
perature between  250°  C.  and  270°  C.  form  acid  salts  of 
the  type  RHSO4  and  the  neutral  salts  of  the  type  R2SO4  on 
ignition. 

Some  tendency  of  the  caesium  salt  to  hold  more  sulphuric 
acid  than  corresponds  to  the  formation  of  the  acid  sulphate 
RHSO4  was  apparent  at  temperatures  between  258°  C.  and 
270°  C.,  but  upon  raising  the  temperature  above  300°  C.  the 
loss  was  excessive  and  showed  a  tendency  on  the  part  of  the 
acid  sulphate  to  pass,  at  this  temperature,  toward  the  con- 
dition of  the  pyrosulphate. 

Sodium  and  potassium  salts,  when  heated  under  the  con- 
ditions described,  give  pyrosulphates  of  the  type  R2S2O7  which 
on  ignition  go  into  the  neutral  sulphate  of  the  form  R2SO4. 
Lithium  gives  neither  salts  of  the  type  RHSO4  nor  R2S2O7 
under  the  conditions  of  these  experiments. 


XLVIII 

THE    ESTIMATION     OF     CALCIUM,    STRONTIUM, 
AND  BARIUM  AS  THE   OXALATES. 

BY  CHARLES  A.  PETERS. 

A  FORMER  article  from  this  laboratory*  describes  the  condi- 
tions under  which  oxalic  acid  may  be  titrated  by  potassium 
permanganate  in  the  presence  of  hydrochloric  acid,  and  states 
that  the  extra  consumption  of  permanganate  which  ordinarily 
takes  place  when  oxalic  acid  is  titrated  by  permanganate  in 
the  presence  of  hydrochloric  acid,  may  be  prevented  by  the 
addition  of  a  manganous  salt.  This  fact  led  to  the  idea  of 
effecting  the  solution  of  the  alkaline  earth  oxalates  in  hydro- 
chloric acid  and  titrating  the  free  oxalic  acid  with  perman- 
ganate in  the  presence  of  a  manganous  salt,  and  so  to  the 
study  of  the  conditions  under  which  precipitates  of  strontium 
and  barium  oxalates  could  be  obtained  sufficiently  insoluble 
for  quantitative  purposes,  the  conditions  under  which  calcium 
oxalate  is  insoluble  being  already  well  known. 

The  permanganate  solution  for  this  work  was  standardized 
against  freshly  recrystallized  ammonium  oxalate,  and  on 
oxalic  acid,  the  standards  agreeing. 

Calcium  Oxalate. 

It  is  well  known  that  calcium  may  be  estimated  by  treating 
the  precipitated  oxalate  with  sulphuric  acid  and  titrating  by 
permanganate  the  oxalic  acid  set  free.f  In  the  work 
described  in  the  present  article,  the  precipitate  of  calcium 
oxalate  has  been  dissolved  in  hydrochloric  acid  and  the  oxalic 

*  Gooch  and  Peters,  Am.  Jour.  Sci.,  vii,  461.    This  volume,  p.  222. 
t  Mohr,  Titrirmethode,  6.  Aufl.,  S.  227. 


374 


ESTIMATION  OF  CALCIUM,   STRONTIUM, 


acid  titrated  by  permanganate  in  the  presence  of  a  manganous 
salt.  The  process  was  as  follows :  The  boiling  hot  solution 
of  calcium  chloride  was  precipitated  with  ammonium  oxalate, 
allowed  to  stand  12  hours,  and  the  supernatant  liquid  de- 
canted on  asbestos.  The  precipitate  was  washed  two  or  three 
tunes  by  decantation  with  50-100  cm3  of  cold  water  and 
brought  on  the  felt.  The  crucible  containing  the  precipitate 
was  returned  to  the  beaker,  100-200  cm3  of  water  were 
added,  together  with  5-10  cm3  of  strong  hydrochloric  acid  and 
0.5-1.0  grin,  of  manganous  chloride,  and  the  oxalic  acid 
titrated  at  a  temperature  of  35° -45°.  The  results,  given  in 
Table  I,  are  obviously  excellent. 

TABLE  I. 


CaO  taken  as 
CaClj. 

Ammonium 
oxalate. 

Volume  at 
precipitation. 

CaO  found. 

Error. 

grm. 

grm. 

cm3 

gnu. 

grm. 

0.0656 

0.3 

100 

0.0657 

0.0001+ 

0.0656 

0.3 

100 

0.0656 

0.0000 

0.0656 

0.3 

150 

0.0658 

0.0002+ 

0.0656 

0.3 

100 

0.0655 

0.0001- 

0.0985 

0.5 

175 

0.0981 

0.0004- 

0.1313 

0.6 

150 

0.1315 

0.0002+ 

0.1313 

0.6 

200 

0.1315 

0.0002+ 

Extended  washing  with  hot  water,  however,  is  to  be  avoided 
after  the  precipitant,  ammonium  oxalate,  has  been  removed. 
In  one  experiment,  for  example,  in  which  the  precipitate,  on 
the  felt,  was  washed  fourteen  times  with  portions  of  about 
50  cm3  each  of  hot  water,  each  portion  bleached  from  2-6 
drops  of  approximately  ^  permanganate,  making  a  total  loss 
of  0.0034  grm.  of  calcium  oxide. 

Strontium  Oxalate. 

Souchay  and  Lenssen  *  state  that  strontium  oxalate  is  solu- 
ble in  12,000  parts  of  water.  This  fact  would  seem  sufficient 
to  warrant  the  study  of  the  quantitative  separation  of  stron- 
tium as  the  oxalate.  In  the  work  which  follows  strontium 


Ann.  Chem.  (Liebig),  cii,  35. 


AND  BARIUM  AS   THE  OXALATE  S. 


375 


oxalate  has  been  precipitated  both  in  alcoholic  solution  and 
in  water  solution,  and  for  convenience  these  two  conditions 
of  precipitation  will  be  discussed  separately. 

All  the  strontium  salts,  of  established  purity,  were  stan- 
dardized by  precipitation  with  sulphuric  acid  in  a  solution 
containing  at  least  one-half  its  volume  of  alcohol,  and  with 
some  solutions  confirmatory  standards  were  also  obtained  by 
evaporation  with  sulphuric  acid. 

Precipitation  in  Alcoholic  Solution.  To  determine  the  com- 
pleteness of  the  precipitation  in  alcoholic  solution  strontium 
nitrate  was  precipitated  by  ammonium  oxalate  in  a  solution 
containing  one-third  of  its  volume  of  alcohol,  the  mixture  was 
allowed  to  stand  over  night,  the  liquid  was  filtered  off  on 
asbestos,  and  the  precipitate  was  treated  in  the  capped  filter- 
ing crucible  with  sulphuric  acid,  ignited,  and  weighed  as  the 
sulphate.  The  results  are  given  in  Table  II.  It  is  plain  from 

TABLE  IL 


SrO  taken  as 
Sr(N03)2. 

Ammonium 
oxalate. 

Volume  at 
precipitation. 

Volume  of 
alcohol. 

SrO  found 

as  SrS04. 

Difference. 

grm. 

grm. 

cm* 

griii. 

griii. 

0.2434 

0.8 

180 

0.2440 

0,0006+ 

0.2434 

0.8 

180 

0.2437 

0.0003+ 

0.0022 

0.2 

100 

0.0022 

0.0000 

0.0013 

0.2 

100 

0.0014 

0.0001+ 

0.0004 

0.04 

100 

0.0004 

0.0000 

the  results  in  this  table  that  the  precipitation  of  even  small 
amounts  of  the  strontium  salt  from  a  solution  containing  one- 
third  of  its  volume  of  alcohol  is  practically  complete. 

To  determine  the  minimum  amount  of  alcohol  necessary  for 
the  complete  precipitation  of  the  strontium  oxalate,  experiments 
were  made  using  varying  proportions  of  85  per  cent  alcohol 
with  different  amounts  of  ammonium  oxalate,  and  the  filtrates 
from  such  experiments  were  tested  for  strontium  by  the  addi- 
tion of  more  alcohol.  The  results  given  in  Table  III  show 
that  when  a  moderate  excess  of  ammonium  oxalate  is  present, 
a  volume  of  85  per  cent  alcohol,  amounting  to  one-fifth  of 


376          ESTIMATION  OF  CALCIUM,   STRONTIUM, 

TABLE  III. 


SrO  present 
as  Sr(N03)3. 

Ammonium 
oxalate. 

Volume  of 
liquid. 

Proportion 
of  85  per  cent 

alcohol. 

SrO  found  in 
filtrates,  weighed 
as  SrSO4. 

grin. 

grm. 

cm3 

grm. 

0.1 

0.4 

100 

1 

0.0000 

0.1 

0.4 

100 

JL 

0.0000 

0.1 

0.4 

100 

1 

0.0004 

0.1 

0.2 

100 

1 

0.0000 

0.1 

0.2 

100 

JL 

0.0009 

0.1 

0.2 

100 

JL 

0.0020 

0.1 

0.1 

100 

* 

0.0002 

the  whole,  is  sufficient  to  complete  the  precipitation  of  the 
strontium  as  the  oxalate. 

The  conditions  under  which  strontium  oxalate  is  insoluble 
having  been  determined,  the  process  for  the  volumetric  esti- 
mation of  strontium  was  carried  out  as  follows:  The  hot 
solution  of  a  strontium  salt  was  precipitated  with  ammonium 
oxalate,  85  per  cent  alcohol,  amounting  to  from  one-fifth  to  one- 
third  the  total  volume,  was  added,  the  mixture  was  allowed  to 
stand  over  night,  and  the  clear  liquid  was  decanted  on  an 
asbestos  filter.  The  precipitate  was  washed  with  a  mixture 
of  equal  parts  of  85  per  cent  alcohol  and  water,  transferred  to 
the  filter,  dried  in  the  filtering  crucible  over  a  flame  to  free 
it  from  alcohol,  returned  to  the  beaker  previously  dried,  treated 
with  sulphuric  acid,  or  with  5-10  cm3  of  hydrochloric  acid 
(in  the  latter  case  0.5-1.0  grm.  of  a  manganous  salt  being 
added)  and  the  liberated  oxalic  acid  was  titrated  by  perman- 
ganate. The  results  obtained  by  this  method  are  accurate 
and  are  given  in  Table  IV. 

In  the  last  experiment  in  which  a  comparatively  large 
amount  of  strontium  salt  was  present  and  the  dilution  low, 
there  is  a  slight  tendency  towards  a  minus  error,  due  probably 
to  the  occlusion  of  some  oxalic  acid  by  the  strontium  sulphate 
formed.  This  phenomenon  would  favor  titration  at  greater 
dilution  when  sulphuric  acid  is  used  to  liberate  the  oxalic 
acid  from  large  amounts  of  strontium  oxalate. 

Precipitation  in  Water  Solution.     In  order  to  determine  the 


AND  BARIUM  AS  THE  OXALATES. 


377 


TABLE  IV. 
VOLUME  DURING  TITBATION  150-250  CMS. 


SrO  taken 
as 

Ammonium 

Volume  at 
precipita- 
tion. 

Propor- 
tion of 
85  per  cent 
alcohol. 

Acid  pres- 
ent during 
titratiou. 

SrO 

found. 

Error. 

grm. 
0.0974 

grm. 
0.4 

cm3 
100 

HC1 

grm. 
0.0973 

grm. 
0.0001- 

0.0974 

0.4 

100 

HC1 

0.0983 

0.0009+ 

0.0974 

0.4 

100 

HC1 

0.0975 

0.0001+ 

0.0974 

0.8 

100 

HC1 

0.0981 

0.0007+ 

0.1948 

0.4 

200 

HC1 

0.1943 

0.0005- 

0.1948 

0.8 

200 

HC1 

0.1942 

0.0006- 

0.0974 

0.4 

100 

H2S04 

0.0970 

0.0004- 

0.0974 

0.4 

100 

H2S04 

0.0977 

0.0003+ 

0.0974 

0.4 

100 

] 

H.2S04 

0.0976 

0.0002+ 

0.1948 

0.6 

150 

H2S04 

0.1938 

0.0010- 

degree  of  precipitation  of  strontium  salts  in  water  solution, 
strontium  oxide,  taken  as  the  nitrate,  was  precipitated  by 
ammonium  oxalate,  the  mixture  was  allowed  to  stand  over 
night,  filtered  on  asbestos,  the  precipitate  was  washed  with 
water  containing  one-half  its  volume  of  85  per  cent  alcohol, 
treated  hi  the  capped  crucible  with  a  few  drops  of  sulphuric 
acid,  ignited,  and  weighed  as  the  sulphate.  The  result  gave 
0.0973  grm.  of  strontium  oxide  instead  of  0.0974  grm.  taken. 
The  precipitation,  therefore,  of  strontium  oxalate,  in  water 
solution  with  a  sufficient  excess  of  ammonium  oxalate  present, 
is  practically  complete. 

To  determine  the  amount  of  ammonium  oxalate  necessary 
for  the  precipitation  of  strontium  salts  in  water  solution, 
experiments  were  made  in  which  strontium  oxalate  was  pre- 
cipitated in  the  presence  of  varying  amounts  of  ammonium 
oxalate,  allowed  to  stand  over  night,  the  clear  liquid  was 
decanted  on  asbestos,  and  the  precipitate  was  washed  twice 
with  10-20  cm3  of  cold  water.  The  results  obtained  by 
the  estimation  of  the  oxalic  acid  by  permanganate  show 
that  an  amount  of  ammonium  oxalate  several  times  larger 
than  that  required  for  the  theoretical  formation  of  strontium 
oxalate  is  necessary  for  the  separation  of  the  strontium  oxalate. 
The  experiments  are  recorded  in  Table  V. 


378          ESTIMATION  OF  CALCIUM,   STRONTIUM, 

TABLE  V. 


SrO, 
taken  as 
Sr(NOa),. 

Ammonium 

oxalate. 

Volume 
at  precipi- 
tation. 

Acid  present 
during 
titratiou. 

SrO 
found. 

Error. 

gnn. 

grm. 

cms 

grm. 

grm. 

0.0487 
0.0487 

0.064 
0.0768 

100 
100 

H2S04 
H2SO4 

0.0441 
0.0465 

0.0046- 
0.0022- 

0.0487 

0.16 

100 

H2SO4 

0.0488 

0.0001+ 

0.0974 

0.128 

100 

H2S04 

0.0939 

0.0025- 

0.0974 

0.16 

100 

H2S04 

0.0959 

0.0015- 

0.0974 

0.32 

100 

H2S04 

0.0976 

0.0001+ 

The  solvent  action  of  a  large  amount  of  water  on  a  pre- 
cipitate of  strontium  oxalate  was  tested  by  washing  a  pre- 
cipitate equivalent  to  0.0974  grm.  of  the  oxide  with  150  cm3 
of  cold  water.  The  precipitate,  when  weighed  as  the  sulphate, 
showed  a  loss  of  0.0033  grm.,  as  the  oxide,  which  amount  was 
subsequently  recovered  from  the  filtrate  by  the  addition  of 
ammonium  oxalate  and  alcohol.  Plainly  excessive  washing 
with  water  is  to  be  avoided.  In  the  estimation,  therefore,  of 
strontium  precipitated  as  the  oxalate  in  water  solution,  the 
amount  of  water  used  in  washing  was  limited.  It  was  found 
that  40-50  cm3  of  water  judiciously  applied  was  sufficient  to 
wash  out  the  ammonium  salt  without  producing  appreciable 
solvent  effect  upon  the  strontium  oxalate.  The  process  of 
treatment  was  similar  to  that  used  in  the  precipitations  from 
alcoholic  solution,  excepting  that  no  alcohol  was  added  to  the 
solution,  that  the  washing  was  effected  with  a  limited  amount 
of  water,  and  that,  there  being  no  alcohol  present  to  effect  the 
titration,  the  precipitate  was  not  dried  before  treatment  with 
permanganate.  The  results  are  given  hi  Table  VI. 

In  the  results  recorded  in  section  A  of  Table  VI,  the  stron- 
tium oxalate  was  treated  with  sulphuric  acid  and  titrated  at 
80°,  the  volume  being  200-300 cm3;  while  hi  the  experiments 
given  in  section  B,  the  precipitate  was  treated  with  hydro- 
chloric acid  and  titrated  at  35°-45°,  at  a  volume  of  100-200  cm3, 
after  the  addition  of  0.5-1.0  grm.  of  manganous  chloride.  The 
results  show  that  0.1  grm.  of  strontium  salt,  calculated  as  the 


AND  BARIUM  AS  THE  OXALATES. 


379 


oxide,  may  be  estimated  as  the  oxalate  with  a  fair  degree  of 
accuracy  when  precipitated  in  100-250  cm3  of  water  by  a 


TABLE  VI. 


SrO, 
taken  as 
Sr(NOs),. 

Ammonium 
oxalate. 

Volume 
at  precipi- 
tation. 

Acid  present 
during 
titration. 

SrO 
found. 

Error. 

A. 

grm. 

0.0974 
0.0974 
0.0974 
0.0974 
0.0974 
0.0974 
0.0974 
0.0974 
0.0974 

grm. 
0.5 
0.5 
0.5 
0.5 
0.8 
0.8 
1.0 
2.0 
2.0 

Cfflg 

100 
100 
100 
100 
100 
100 
100 
100 
100 

H2SO4 
H2S04 
H2S04 
H2S04 
H2S04 
H2S04 
H2S04 
H2S04 
H2SO4 

grm. 
0.0966 
0.0985 
0.0977 
0.0963 
0.0981 
0.0966 
0.0965 
0.0963 
0.0970 

grm. 
0.0008- 
0.0011+ 
0.0003+ 
0.0011- 
0.0007+ 
0.0008- 
0.0009- 
0.0011- 
0.0004- 

SrO, 
taken  as                                                     B. 
8rCl2. 

0.0778 
0.0778 
0.0778 
0.0778 

0.5 
0.5 
0.5 
0.5 

100 
100 
100 
100 

H2S04 
H2S04 
H2SO4 
H2S04 

0.0792 
0.0767 
0.0776 
0.0776 

0.0014+ 
0.0011- 
0.0002- 
0.0002- 

SrO, 
taken  as 
Sr(N08)2. 

0.0974 
0.0974 
0.0974 
0.0974 
0.0974 
0.0974 
0.0974 
0.0974 

0.8 
2.0 
0.8 
0.8 
0.8 
0.8 
0.8 
0.8 

250 
250 
100 
100 
100 
100 
100 
100 

H2S04 
H2S04 
HCI 
HCI 
HCI 
HCI 
HCI 
HCI 

0.0973 
0.0975 
0.0971 
0.0980 
0.0975 
0.0980 
0.0973 
0.0978 

0.0001- 
0.0001+ 
0.0003- 
0.0006+ 
0.0001+ 
0.0006+ 
0.0001- 
0.0004+ 

C. 

0.2425 
0.2436 
0.2436 
0.2436 
0.2436 
0.2436 

0.384 
0.384 
0.64 
0.8 
2.0 
2.0 

125 
125 
125 
125 
125 
125 

H2S04 
H2S04 
H2S04 
H2S04 
H2S04 
H2S04 

0.2376 
0.2402 
0.2411 
0.2367 
0.2376 
0.2402 

0.0049- 
0.0034- 
0.0025- 
0.0069- 
0.0060- 
0.0034- 

D. 

0.2436 
0.2436 
0.2436 
0.2436 

0.8 
0.8 
2.0 
2.0 

250 
250 
250 
250 

H2SO4 
H2S04 
H2S04 
H2S04 

0.2443 
0.2446 
0.2440 
0.2431 

0.0007+ 
0.0010+ 
0.0004+ 
0.0005- 

380          ESTIMATION  OF  CALCIUM,   STRONTIUM, 
TABLE  VI  (continued). 


SrO, 
taken  as 
Sr(N08),. 

Ammonium 

oxalate. 

Volume 
at  precipi- 
tation. 

Acid  present 
during 
titratiou. 

SrO 
found. 

Error. 

E. 

grui. 

grm. 

cm8 

grm. 

grm. 

0.2436 

0.8 

500 

H2S04 

0.2396 

0.0040- 

0.2436 

2.0 

500 

H2S04 

0.2403 

0.0033- 

0.2436 

2.0 

500 

H2S04 

0.2413 

0.0023- 

0.2436 

4.0 

600 

H2S04 

0.2410 

0.0026- 

0.2436 

8.0 

500 

H2S04 

0.2407 

0.0029- 

0.4872 

2.0 

500 

H2SO4 

0.4837 

0.0035- 

0.4872 

4.0 

500 

H2S04 

0.4855 

0.0017- 

0.5430 

5.0 

500 

HnS04 

0.5422 

0.0008- 

0.4579 

10.0 

500 

H2SO4 

0.4554 

0.0025- 

0.7307 

4.0 

500 

HC1 

0.7262 

0.0045- 

sufficient  excess  of  ammonium  oxalate.  In  the  experiments 
recorded  in  section  C,  in  which  the  amount  of  strontium  salt 
in  125  cm3  of  water  is  increased,  a  negative  error  is  intro- 
duced, which  is  not  diminished  by  the  presence  of  a  large 
amount  of  ammonium  oxalate,  but  when  the  dilution  is  in- 
creased to  250  cm3,  as  is  the  case  in  experiments  given  in 
section  D,  so  that  the  conditions  correspond  more  nearly  to 
those  recorded  in  sections  A  and  B,  the  errors  fall  to  a 
minimum.  In  the  experiments  recorded  in  section  E,  in 
which  the  dilution  is  increased  to  500  cm3,  an  error  is  in- 
troduced which  is  not  prevented  by  the  presence  of  a  large 
excess  of  ammonium  oxalate  and  which  is  independent  of  the 
amounts  of  strontium  salt  used. 

Eight  of  the  water  filtrates  and  wash  waters  obtained  in 
the  experiments  recorded  in  Table  VI  were  tested  for  traces 
of  strontium  by  the  addition  of  alcohol,  and  in  all  cases  a 
small  amount  of  strontium  was  found,  amounting,  in  the 
average,  to  0.0010  grm.  in  100  cm3  of  water. 

Barium  Oxalate. 

Barium  oxalate  according  to  Souchay  and  Lenssen  *  is 
soluble  in  2590  parts  of  cold  water,  and  according  to  Berg- 

*  Ann.  Chem.  (Liebig),  xc,  102. 


AND  BARIUM  AS   THE   OXALATES. 


381 


man  *  is  scarcely  at  all  soluble  in  alcohol.  The  attempt  was 
made  to  estimate  barium  by  precipitation  with  ammonium 
oxalate  in  a  mixture  containing  alcohol.  It  was  found  that 
in  filtrates  from  oxalate  precipitations  in  which  0.1-0.2  grm. 
of  barium  oxide,  taken  as  the  nitrate,  had  been  precipitated  in 
volumes  of  100  cm8,  containing  30  cm3  of  absolute  alcohol, 
and  allowed  to  stand  over  night,  treatment  with  sulphuric 
acid  gave  barium  sulphate  amounting  in  the  average  to  no 
more  than  0.0001  grm.  of  barium  oxide.  The  insolubility  of 
barium  oxalate  under  these  conditions,  therefore,  is  practically 
complete. 

The  process  for  the  estimation  of  barium  was  as  follows : 
Ammonium  oxalate  was  added  to  a  solution  of  a  barium  salt, 

TABLE  VII. 


BaO  taken 
as  Ba(N03)2. 

Ammonium 
oxalate. 

Volume  at 
precipita- 
tion. 

Acid  present 
during 
titration. 

BaO 

found. 

Error. 

A. 

grm. 

0.1165 
0.1165 
0.1165 
0.1165 
0.1165 
0.1165 
0.1165 
0.2330 
0.2330 
0.2330 

grin. 

0.2 
0.2 
0.2 
0.2 
0.2 
0.2 
0.2 
0.4 
0.4 
0.4 

cm3 
100 
100 
100 
100 
100 
100 
100 
100 
100 
100 

HCI 
HCI 
HCI 
HCI 
HCI 
HCI 
HCI 
HCI 
HCI 
HCI 

gnu. 

0.1177 
0.1170 
0.1164 
0.1151 
0.1165 
0.1176 
0.1164 
0.2319 
0.2335 
0.2342 

grm. 
0.0012+ 
0.0005+ 
0.0001- 
0.0014- 
0.0000 
0.0011+ 
0.0001- 
0.0011- 
0.0005+ 
0.0012+ 

BaO  taken 
as  BaCl2. 

0.4 

0.4 
0.4 
0.4 
0.4 

100 
100 
100 
100 
100 

HCI 
HCI 
HCi 
HCI 
HCI 

0.0952 
0.0939 
0.0941 
0.1893 
0.1892 

0.0010+ 
0.0003- 
0.0001- 
0.0009+ 
0.0008+ 

0.0942 
0.0942 
0.0942 
0.1884 
0.1884 

B. 

0.0942 
0.1884 
0.0942 

0.2 
0.4 
0.2 

200 
200 
500 

H2S04 
H2S04 
H2S04 

0.0858 
0.1732 
0.0857 

0.0084- 
0.0152- 
0.0085- 

*  Bergman's  Essays,  i,  320. 


382          ESTIMATION  OF  CALCIUM,   STRONTIUM, 

containing  30  per  cent  of  its  volume  of  absolute  alcohol,  the 
mixture  was  allowed  to  stand  over  night,  filtered  on  asbestos, 
the  precipitate  was  washed  by  decantation  with  100-200  cm3 
of  water  containing  30  per  cent  of  absolute  alcohol,  and  dried 
over  a  flame  to  insure  the  removal  of  the  alcohol.  The  cruci- 
ble containing  the  precipitate  was  returned  to  the  beaker  also 
previously  dried  over  a  flame,  100-200  cm3  of  water,  5-10 
cm3  of  strong  hydrochloric  acid,  and  0.5-1.0  grm.  of  manganous 
chloride  were  added,  and  the  solution  was  titrated  at  35° -45° 
with  permanganate.  The  results  of  the  experiments,  given  in 
Table  VII,  A,  show  that  barium,  either  as  the  nitrate  or 
chloride,  may  be  estimated  in  the  manner  described  with  a  fair 
degree  of  accuracy. 

In  the  experiments  given  in  section  B  of  Table  VII,  the 
precipitate  of  barium  oxalate  was  treated  with  sulphuric  acid 
after  the  addition  of  the  stated  amount  of  water.  The  results 
show  a  large  loss  of  oxalic  acid  probably  due  to  the  occlusion 
of  some  of  the  oxalic  acid  by  the  barium  sulphate.  This  fact 
must  prevent  the  use  of  sulphuric  acid  in  an  analytical 
process  which  depends  upon  the  liberation  of  oxalic  acid  from 
barium  oxalate. 

Gravimetric  Estimation  of  the  Oxalates  of  Strontium 
and  Barium. 

It  is  well  known  that  calcium  may  be  weighed  as  the 
carbonate  after  a  careful  ignition  of  the  oxalate,  and  it  would 
seem  probable  that  strontium  might  also  be  weighed  as  the 
carbonate.  Precipitates  of  strontium  oxalate,  on  asbestos, 
were  ignited  in  a  capped  crucible  from  2-8  minutes  in  the 
flame  of  a  Bun  sen  burner  and  weighed  as  the  carbonate,  and 
in  a  single  case  the  carbonate  thus  produced  was  converted  by 
treatment  with  sulphuric  acid  to  the  sulphate  and  weighed  as 
such.  The  results  are  given  in  Table  VIII,  and  while  they 
all  show  a  very  slight  loss,  which  amounts  in  experiment  (3) 
to  one  milligram,  when  one-fourth  of  a  gram  of  strontium 
oxide  taken  as  the  nitrate  was  used,  still  the  results  are  fairly 
accurate. 


AND  BARIUM  AS  THE  OXALATES. 
TABLE  VUL 


383 


Ezp. 

BrO  taken 
as  Sr(NOs),. 

SrO  calculated 
from  SrC03 
found. 

SrO  calculated 
from  SrSO4 
found. 

grm. 

grm. 

grm. 

(1) 

(2) 

0.1120 
0.1120 

0.1113 
0.1116 

.... 

(3) 

0.2435 

0.2425 

0.2437 

Precipitates  of  barium  oxalate  were  also  ignited  from  5-10 
minutes  and  weighed  as  the  carbonate.  The  results  are  given 
in  Table  IX,  and  are  fairly  accurate. 


TABLE  IX. 


Exp. 

BaO  taken 
as  Ba(N03)2. 

BaO  calculated 
from  BaCO8 
found. 

Difference. 

(1) 
(2) 
(3) 

griii. 

0.2912 
0.2912 
0.2912 

grm. 
0.2909 
0.2901 
0.2901 

grm. 
0.0003- 
0.0011- 
0.0011- 

The  results  of  this  work  may  be  summarized  as  follows: 
In  the  estimation  of  calcium  by  titration  of  the  oxalate  with 
permanganate,  accurate  results  may  be  obtained  when  hydro- 
chloric acid  (with  a  manganous  salt)  is  used  as  the  solvent. 
Strontium  salts  may  be  precipitated  by  ammonium  oxalate 
with  practical  completeness  in  a  solution  containing  one-fifth 
of  its  volume  of  85  per  cent  alcohol,  and  with  approximate 
completeness  from  water  solutions  at  a  dilution  not  exceeding 
250  cm3.  Furthermore  strontium  oxalate  may  be  titrated  by 
permanganate  with  accuracy  when  either  sulphuric  acid  or 
hydrochloric  acid  (with  a  manganous  salt)  is  used  to  liberate 
the  oxalic  acid.  Barium  may  be  precipitated  with  practical 
completeness  by  ammonium  oxalate  in  a  solution  containing 
30  per  cent  of  alcohol,  and  the  barium  oxalate  thus  obtained 
may  be  dissolved  in  hydrochloric  acid  and  titrated  by  per- 
manganate after  the  addition  of  a  manganous  salt.  Strontium 
and  barium  oxalates  may  be  converted  to  carbonates  by 
ignition,  and  weighed  as  such. 


XLIX 

THE  ACTION  OF  SODIUM  THIOSULPHATE  ON 
SOLUTIONS  OF  METALLIC  SALTS  AT  HIGH 
TEMPERATURES  AND  PRESSURES. 

BY  JOHN  T.  NORTON,  JR. 

THE  use  of  sodium  thiosulphate  as  a  substitute  for  hydro- 
gen sulphide  in  effecting  precipitations  and  its  application 
in  the  case  of  arsenic,  antimony,  copper,  and  platinum  was 
suggested  by  Himly  *  before  the  middle  of  the  present  cen- 
tury. Thirteen  years  later  Vohl  f  and  Slater,  independently  $ 
drew  attention  to  this  use  of  sodium  thiosulphate  and  ex- 
tended the  investigation  to  salts  of  tin,  mercury,  silver,  gold, 
lead,  bismuth,  and  cadmium.  Slater  in  addition  studied  the 
action  of  sodium  thiosulphate  upon  chromic  acid,  molybdates, 
ferrous  and  ferric  ferrocyanides,  ferric  sulphocyanides  and 
potassium  permanganate.  Following  out  these  lines,  the 
precipitation  of  copper,  together  with  arsenic  antimony,  by 
treating  with  sodium  thiosulphate  the  hot  solution  contain- 
ing sulphuric  acid,  and  the  separation  of  these  elements  from 
tin,  zinc,  iron,  nickel,  cobalt,  and  manganese  has  been  advo- 
cated by  Westmoreland  ;§  and  quite  recently  Faktor  ||  has 
studied  the  action  of  sodium  thiosulphate  upon  neutral  salts 
of  several  of  the  elements  mentioned,  as  well  as  the  modi- 
fying influence  of  ammonium  chloride  and  other  salts  upon 
the  course  of  the  reaction. 

Subsequently  to  the  work  of  Himly,  Vohl,   and  Slater, 
Chancel  **  developed  his  well  known  method  for  the  precipi- 

*  Ann.  Chem.  (Liebig),  xliii,  150. 
t  Ann.  Chem.  (Liebig),  xcvi,  237. 

t  Chemical  Gazette,  1855,  p.  369.  §  Jour.  Soc  Chem.  Ind.,  v,  61. 

||  Chem.  Centralblatt,  1900,  ii,  20,  67,  239,  594. 
»*  Comp.  rend.,  xlvi,  987. 


ACTION  OF  SODIUM   THIOSULPHATE,  ETC.       385 

tation  of  aluminum  as  the  hydroxide  and  its  separation  from 
salts  of  iron  by  boiling  with  sodium  thiosulphate  the  nearly 
neutral  solution,  containing  the  salt  of  aluminum  and  iron, 
at  suitable  dilution ;  and  upon  an  extension  of  the  principle 
of  Chancel's  separation  of  aluminum  from  iron  Stromeyer  * 
founded  his  well  known  processes  for  the  separation  of  titan- 
ium and  zirconium  from  iron.  The  latter  process  appears 
to  be  fairly  trustworthy ;  but  of  Chancel's  method,  although 
it  has  met  with  wide  acceptance,  it  was  shown  by  Wolcott 
Gibbs,  very  soon  after  its  announcement,!  that  it  fails  to  bring 
about  complete  separation  of  alumina  within  a  reasonable 
period  of  boiling,  and  this  result  has  been  confirmed  by  Zimmer- 
man, J  who  has  shown  that  the  boiling  must  be  continued  fifteen 
hours  in  order  to  complete  the  precipitation  of  the  alumina. 

It  was  shown  by  Dr.  Gibbs  that  when  the  treatment  of  salts 
of  aluminum  by  thiosulphate  was  carried  on  in  sealed  tubes 
under  pressure  at  120°  C.,  the  precipitation  of  alumina  was 
complete,  and  further  that  the  precipitation  of  sulphides  of 
nickel,  cobalt,  and  iron,  though  partial  under  ordinary  atmos- 
pheric pressure,  was  made  complete  by  heating  in  sealed  tubes 
to  120°-140°  C. 

In  repeating  the  experiments  of  Dr.  Gibbs  qualitatively  and 
extending  them,  I  have  made  use  of  the  well  known  Pfungst 
tube  to  secure  the  necessary  pressure.  In  each  experiment 
a  test-tube  containing  the  mixture  of  an  excess  of  sodium 
thiosulphate  with  the  salt  whose  action  was  studied  was 
placed  within  the  Pfungst  tube  containing  some  water,  the 
cover  of  the  latter  was  set  in  place  and  firmly  bolted  upon 
a  washer  of  lead,  and  the  whole  was  submitted  to  tempera- 
tures varying  from  140°  to  200°  C.  for  an  hour  by  immersing 
in  a  bath  of  paramne.  After  cooling,  the  test-tube  was  taken 
out,  the  precipitate  was  filtered  off,  and  the  filtrate  tested  by 
appropriate  reagents  to  determine  the  completeness  of  pre- 
cipitation. The  following  table  records  the  details  of  these 
experiments. 

*  Ann.  Chem.  (Liebig),  cxiii,  127. 

t  Zeit.  anal.  Chem.,  iii,  389.  t  Inaug.  Diss.  Berlin,  1887. 

VOL.  u.  —  25 


386 


ACTION  OF  SODIUM  THIOSULPHATE  ON 


TABLE  I. 
ACTION  OP  NA2S2O3  ON  SALTS  UNDER  PRESSURE. 


Salts  used. 

Precipitates. 

Degree  of 
precipitation. 

SULPHIDES. 

NiSO4. 
CoSO4. 
FeCl3. 
ZnSO4. 
Pb02(C2H80)a. 
Hg(N03V 
AgN08. 
CuSO4. 
CdS04. 
KSbC4H4(X. 
Bi(N08)8. 

NiS  +  S. 
CoS  +  S. 
FeS  +  S. 
ZnS  +  S. 
PbS  +  S. 
HgS  +  S. 
Ag2S  +  S. 
CuS,Cu2S,  +  S. 
CdS  +  S. 
Sb2S8  +  S. 
Bi2S8  +  S. 

Complete. 
Complete. 
Complete. 
Complete. 
Complete. 
Complete. 
Complete. 
Complete. 
Complete. 
Complete. 
Complete. 

HYDROXIDES. 

(NH4)A1(S04)2.12H20. 
K2Cr2O7. 
K2ZrF6. 
KoTiFa. 
Th(N08)4. 

A108H3  +  S. 
Cr03H3  +  S. 
Zr04H4  +  S. 
Ti04H4  +  S. 
Th04H4  +  S. 

Complete. 
Complete. 
Complete. 
Complete. 
Complete. 

ELEMENTS. 

SeO2. 
Te02. 

Se  +  S. 
Te  +  S. 

Complete. 
Complete. 

SULPHIDES. 

MnS04. 
AuCl8. 
(NH4)2Mo04. 

MnS  +  S. 
Au2S  +  S. 
MoS8(?)  +  S  —  Red  liquid. 

Partial. 
Partial. 
Partial. 

HYDROXIDES. 

BeCLj. 

Be02H2  +  S. 

Partial. 

UNDETERMINED. 

(NH4)2U207. 
K2PtCle. 
CeCl8. 
CaCl2. 
SrCL. 
Bad-. 
MgS04. 
NH4VO8. 
H2KAsO4. 

Black. 
Gray,  reddish  brown  liquid. 
White,  yellow  liquid. 
White,  yellow  liquid. 
White,  yellow  liquid. 
White,  yellow  liquid. 

,  Brown  liquid. 

Partial. 
Partial. 
Partial. 
Partial. 
Partial. 
Partial. 
None. 
None. 
None. 

SOLUTIONS  OF  METALLIC  SALTS.  387 

A  perusal  of  this  table  brings  to  light  several  interesting 
facts.  It  appears  that  salts  of  nickel,  cobalt,  iron,  zinc,  lead, 
mercury,  silver,  copper,  cadmium,  antimony  and  bismuth  are 
completely  precipitated  as  sulphides  by  sodium  thiosulphate 
under  the  prevailing  conditions  of  temperature  and  pressure. 
In  the  case  of  manganese  precipitation  is  only  partial,  and 
arsenic  does  not  seem  to  be  precipitated  from  an  arsenate  with- 
out the  addition  of  acid.  Tin,  curiously  enough,  is  not  thrown 
down  as  the  sulphide  from  a  stannous  salt,  but  gives  a  dirty 
white  precipitate  of  uncertain  composition.  Salts  of  aluminum, 
chromium,  titanium,  zirconium  and  thorium  are  completely 
precipitated  as  the  hydroxides ;  but  in  the  case  of  beryllium, 
which  one  might  expect  to  act  similarly,  the  precipitation  as 
the  hydroxide  is  incomplete.  Salts  of  selenium  and  tellurium 
are  reduced,  and  the  elements  are  precipitated.  The  precipi- 
tates obtained  with  barium,  strontium,  and  calcium  were 
white  in  a  bright  yellow  liquid,  but  no  study  was  made  of  the 
constitution  of  either  precipitate  or  liquid.  In  the  case  of 
magnesium  there  was  no  precipitate.  Salts  of  molybdenum, 
vanadium  and  uranium  gave  dark  colored  liquids.  Thallium 
yielded  a  white  spongy  mass  which  on  compression  was  re- 
duced to  a  very  small  bulk  without  disintegrating.  Salts  of 
gold  and  platinum  gave  slight  dark  precipitates,  presumably 
sulphides,  surrounded  by  dark  colored  liquids. 

The  apparatus  used  in  these  experiments  and  described 
above  is  easily  handled  and  answers  sufficiently  well  for 
qualitative  purposes.  But,  obviously,  the  introduction  into 
precipitates  of  foreign  matter  caused  by  the  action  of  water 
on  the  glass  of  the  test-tube  and  porcelain  lining  of  the 
Pfungst  tube,  precludes  the  possibility  of  an  exact  quanti- 
tative study  of  the  reactions  involved.  For  the  subsequent 
experiments,  therefore,  conducted  upon  the  same  general 
lines,  a  digester  with  an  interior  cylindrical  cavity  of  about 
12  cm.  in  depth  by  5  cm.  in  diameter,  and  provided  with  a 
pressure  gauge  was  employed.  As  a  container  for  the  solu- 
tions to  be  tested,  use  was  made  of  a  platinum  cylinder,  4  cm. 
in  diameter  and  10  cm.  deep,  provided  with  a  loose  cover. 


388 


ACTION  OF  SODIUM  THIOSULPHATE   ON 


With  this  apparatus  the  following  quantitative  experiments 
which  deal  with  those  elements  which  are  precipitated  as 
hydroxides  —  namely,  aluminum,  beryllium,  chromium,  zir- 
conium, and  titanium  —  were  made. 

In  each  case  a  weighed  quantity  of  the  salt  taken  for  the 
experiment  was  dissolved  in  50  cm3  of  water  in  the  platinum 
vessel  and  to  this  a  known  amount  of  sodium  thiosulphate 
was  added.  The  vessel  was  placed  in  the  digester,  and  the 
latter  was  heated  by  a  Bunsen  burner  in  the  customary  way 
until  the  required  pressure  was  shown  on  the  gauge.  The 
apparatus  was  then  cooled  and  the  platinum  vessel  removed 
from  the  digester.  The  precipitate  was  filtered  off  on  ashless 
paper,  ignited,  and  weighed. 

Experiments  with  a  Salt  of  Aluminum. 

In  a  series  of  experiments  made  according  to  the  method  of 
Chancel,  the  results  of  which  are  shown  in  Table  II,  the 
solution  in  water  of  a  weighed  portion  of  pure  ammonium 
alum  was  treated  with  an  excess  of  sodium  thiosulphate  and 
boiled  vigorously  for  periods  varying  from  ten  minutes  to 
half  an  hour. 

TABLE  II. 


Ezp. 

Amount  of 
Alum  taken 

Amount  of 

A1203 
found. 

Error. 

as  A1,O3. 

grm. 

grm. 

grm. 

grm. 

(1) 

0.0537 

Large  excess. 

0.0471 

0.0066- 

(2) 

0.0537 

Large  excess. 

0.0397 

0.0140- 

is 

0.1083 

Large  excess. 

0.0931 

0.0152- 

(4) 

0.1137 

5  grin. 

0.0979 

0.0158- 

(5) 

0.1139 

2  grm. 

0.1002 

0.0137- 

The  results  substantiate  the  observations  of  Gibbs  *  and  of 
Zimmerman  f  and  show  clearly  that  the  boiling  of  solutions  of 
the  aluminum  salt  and  sodium  thiosulphate  for  a  reasonable 
time  does  not  effect  the  complete  precipitation  of  aluminum  as 
the  hydroxide. 

Table  III   shows    the  result  of    submitting  solutions  of 

*  Loc.  cit.  t  Loc.  cit. 


SOLUTIONS  OF  METALLIC  SALTS. 


389 


ammonium  alum  treated  with  varying  quantities  of  sodium 
thiosulphate  to  a  pressure  of  20  atmospheres  in  the  digester. 
It  usually  required  about  40  minutes  to  raise  the  pressure  to 
the  limit  set;  but  this  limit  once  reached,  the  digester  was 
allowed  to  cool  slowly.  The  duration  of  an  experiment  was 
about  two  hours. 

TABLE  III. 


Alum  taken 
as  A1203. 

Amount  of 
Na,S203  used. 

A1303  found. 

Error. 

grm. 

grm. 

grm. 

grm. 

0.0565 

5 

0.0633 

0.0068+ 

0.1132 

10 

0.1154 

0.0022-f 

0.1153 

5 

0.1186 

0.0033-f- 

0.1128 

3 

0.1129 

0.0001-f- 

0.1126 

3 

0.1142 

0.0016-f- 

0.1128 

2 

0.1120 

0.0008- 

0.1136 

2 

0.1121 

0.0015— 

0.1128 

2.5 

0.1136 

0.0008-f- 

0.1124 

2.5 

0.1127 

0.0003+ 

0.1134 

2.25 

0.1133 

0.0001- 

This  table  shows  that  sodium  thiosulphate  precipitates 
aluminum  completely  as  the  hydroxide  when  pressure  is 
employed.  The  high  results  seen  in  some  of  the  experiments 
appear  to  be  due  to  the  difficulty  of  removing  by  ignition  the 
large  amounts  of  sulphur  found  in  the  action,  as  well  as  to  the 
salts  mechanically  included  in  the  precipitate.  The  amounts 
of  sulphur  and  contaminating  salts  present  depend  upon  the 
amount  of  thiosulphate  taken ;  therefore  this  should  be  as 
small  as  possible,  2-3  grm.  being  sufficient  to  precipitate 
all  the  alumina  in  a  gram  of  alum.  When  the  amount  of 
thiosulphate  is  reasonably  restricted  the  weights  of  alumina 
obtained  accord  fairly  well  with  the  theory. 

Experiments  with  a  Salt  of  Chromium. 

Up  to  the  tune  of  the  completion  of  this  work  nothing 
appears  to  have  been  done  upon  the  precipitation  of  chromium 
as  the  hydroxide  by  means  of  sodium  thiosulphate.  Slater  * 
and  Rose  f  make  mention  of  the  action  of  sodium  thiosulphate 

*  Loc.  cit.  t  Traite  de  Chimie  Analytique,  vol.  i,  p.  479. 


390 


ACTION  OF  SODIUM  THIOSULPHATE  ON 


upon  chromic  acid,  bichromates,  and  neutral  chromates,  but 
give  no  quantitative  data.  Recently,  however,  F.  Faktor  *  has 
studied  the  action  of  sodium  thiosulphate  on  chromium 
compounds.  This  investigator  has  found  that  if  aqueous 
solutions  of  potassium  bichromate  and  sodium  thiosulphate 
are  boiled  together  a  brown  precipitate  of  hydrated  Cr2O8, 
CrO8  separates  out  and  the  liquid  turns  yellow  owing  to  the 
formation  of  normal  chromate.  A  solution  of  potassium 
chromate  is  unaffected  by  boiling  with  thiosulphate  but  in 
presence  of  ammonium  or  of  magnesium  chloride  the  chromium 
is  separated  rapidly  and  completely  in  the  same  form  as  with 
the  bichromate,  and  after  continued  boiling  with  an  excess  of 
thiosulphate  all  the  chromium  present  is  precipitated.  Faktor 
also  found  that  a  solution  of  chromic  chloride  is  completely 
decomposed  by  continued  boiling  with  thiosulphate,  chromic 
hydroxide  and  sulphur  being  precipitated. 

In  the  experiments  shown  in  Table  IV  a  weighed  quantity 
of  pure  potassium  bichromate  was  dissolved  in  water,  a  known 
amount  of  sodium  thiosulphate  added,  and  the  whole  submitted 
to  a  pressure  of  20  atmospheres  in  the  digester.  After  cooling, 
the  precipitate  was  filtered  off  on  an  ashless  paper,  ignited  and 
weighed  as  Cr2O3. 

TABLE  IV. 


Exp. 

K,Cr2O7  taken 
as  Cr2O3. 

Amount  of 
Na2S,06. 

Cr2Os  found. 

Error. 

grm. 

grm. 

grm. 

grm. 

(1) 

0.1330 

3 

0.1341 

0.0011+ 

(2) 

0.1330 

2.5 

0.1326 

0.0004- 

(3) 

0.1322 

2.5 

0.1318 

0.0004- 

(4) 

0.1303 

2 

0.1303 

0.0000 

(5) 

0.1301 

2 

0.1310 

0.0009+ 

6) 

0.1320 

2 

0.1322 

0.0002+ 

The  results  of  these  experiments  are  very  satisfactory,  and 
show  that  under  pressure  sodium  thiosulphate  precipitates 
chromium  rapidly  and  completely  as  the  hydroxide.  It  is 
advisable  to  use  as  small  a  quantity  of  thiosulphate  as  possible 


Zeitschr.  anal.  Chem.,  1900,  xxxix,  345. 


SOLUTIONS  OF  METALLIC  SALTS.  391 

in  order  to  prevent  the  presence  of  much  free  sulphur  in  the 
precipitate. 

Experiments  with  a  Salt  of  Beryllium. 

In  experiments  dealing  with  beryllium  the  salt  used  was  the 
chloride,  a  certain  amount  of  which  was  dissolved  in  water 
diluted  to  a  liter  and  the  amount  of  beryllium  present  deter- 
mined by  precipitating  with  ammonia  and  weighing  as  the 
oxide.  Measured  quantities  of  this  solution  were  drawn  from 
a  burette  as  required.  "When  a  solution  of  a  salt  of  beryllium 
and  sodium  thiosulphate  are  merely  boiled  together  nearly  all 
the  beryllium  remains  in  solution.  It  was  expected  that  the 
use  of  pressure  Would  'throw  out  all  the  beryllium,  but, 
curiously  enough,  when  solutions  of  beryllium  chloride  and 
sodium  thiosulphate  were  submitted  in  the  digester  to  pressures 
ranging  from  10  to  80  atmospheres  only  a  partial  precipitation 
of  the  hydroxide  took  place. 

Experiments  with  Salts  of  Zirconium. 

To  prepare  a  standard  solution  of  the  salt  of  zirconium  it 
was  found  to  be  most  convenient  to  heat  the  double  fluoride 
of  potassium  and  zirconium  with  sulphuric  acid,  evaporate  to 
dryness  in  platinum,  dissolve  the  zirconium  sulphate  remaining 
in  water  and  enough  sulphuric  acid  to  prevent  the  precipitation 
of  the  basic  salt,  and  dilute  to  standard  volume.  Measured 
portions  of  the  solution  were  taken  from  a  burette  as  required 
for  the  experiments.  The  presence,  however,  of  so  large  an 
amount  of  sulphuric  acid  as  was  necessary  to  keep  the 
zirconium  salt  in  solution  tends  to  decompose  sodium  thio- 
sulphate so  rapidly  that  it  was  found  necessary  to  nearly 
neutralize  the  solution  with  ammonium  carbonate  before 
adding  the  sodium  thiosulphate.  The  solution  of  zirconium 
sulphate  was  standardized  by  precipitating  with  ammonia  and 
weighing  as  the  oxide. 

In  experiment  (1)  of  Table  V,  the  solutions  of  zirconium 
sulphate  and  sodium  thiosulphate  were  boiled  together  for  a 
few  minutes  and  then  the  precipitate  filtered  off,  ignited,  and 


392 


ACTION  OF  SODIUM  THIOSULPHATE  ON 


weighed  as  the  oxide.  In  experiments  (2)-(5)  inclusive  similar 
solutions  were  submitted  to  a  pressure  of  20  atmospheres  in 
the  digester. 


TABLE  V. 


Exp. 

Zr02 
taken. 

Na2S203 
taken. 

ZrO. 
found. 

Error. 

grin. 

grru. 

grm. 

grm. 

(1) 

0.0658 

3 

0.0651 

0.0007- 

2 

0.0658 

3 

0.0676 

0.0016+ 

(3) 

0.0666 

2 

0.0670 

0.0004+ 

(4) 

0.0641 

2 

0.0648 

0.0007+ 

(5) 

0.0641 

2 

0.0645 

0.0004+ 

These  results  clearly  show  that  sodium' thiosulphate  precipi- 
tates zirconium  completely  as  the  hydroxide  either  with  or 
without  the  aid  of  pressure. 

Experiments  with  a  Salt  of  Titanium. 

The  solution  of  the  salt  of  titanium  was  obtained  by  treating 
the  double  fluoride  of  potassium  and  titanium  with  sulphuric 
acid,  evaporating  to  dryness,  and  dissolving  the  residue  in 
sulphuric  acid  and  water.  The  solution  was  standardized  by 
precipitating  the  titanium  hydroxide  with  ammonia  and  then 
adding  an  excess  of  acetic  acid  as  recommended  by  Gooch.* 
This  method  of  procedure  avoids  the  tendency  to  excessive 
weight  observed  when  the  titanium  hydroxide  is  precipitated 
by  ammonia  in  presence  of  salts  of  the  alkalies. 

In  the  following  table  is  shown  the  effect  of  treating  a 
solution  of  titanium  sulphate  with  sodium  thiosulphate. 

TABLE  VI. 


Exp. 

Ti02 
taken. 

Na2S20 
taken. 

TiO2 

found. 

Error. 

grm. 

grm. 

grm. 

gnn. 

(1) 

0.0240 

2 

0.0237 

0.0003- 

(2) 

0.0240 

2 

0.0240 

0.0000 

(3) 

0.0240 

2 

0.0240 

0.0000 

*  Am.  Chem.  Jour.,  vii,  285. 


SOLUTIONS  OF  METALLIC  SALTS.  393 

Experiment  (1)  was  conducted  by  merely  boiling  a  solution  of 
the  reagents  named  above,  filtering  off  the  precipitate  and 
weighing  as  the  oxide.  In  experiments  (2)  and  (3)  the 
solution  of  titanium  sulphate  and  sodium  thiosulphate  was 
submitted  to  a  pressure  of  20  atmospheres  in  the  digester. 

These  results  show  that  titanium  is  completely  precipitated 
by  sodium  thiosulphate  either  with  or  without  the  aid  of 
pressure. 

To  recapitulate :  —  I  have  shown  that  sodium  thiosulphate 
will  completely  precipitate  aluminum,  chromium,  zirconium 
and  titanium  as  the  hydroxides  with  the  aid  of  high  tempera- 
ture and  pressure.  Beryllium  is  only  partially  precipitated 
under  similar  conditions.  Mere  boiling  for  a  reasonable  time 
will  not  precipitate  aluminum  and  chromium,  but  it  is  suffi- 
cient in  the  case  of  zirconium  and  titanium. 


SYSTEMATIC  INDEX. 


LABORATORY  APPLIANCES  AND  PREPARATIONS. 

Laboratory  Apparatus  (Gooch),  I,  141.  —  Generation  of  Chlorine  (Gooch  and 
Kreider),  1, 260.  —  Preparation  of  Perchloric  Acid  (Kreider),  1, 282.  —  Labora- 
tory Apparatus  (Kreider),  I,  306. 

INORGANIC  CHEMISTRY. 

Interaction  of  Potassium  Permanganate  and  Sulphuric  Acid  (Gooch  and  Dan- 
ner),  I,  145.  —  Reducing  agents  on  lodic  Acid  (Roberts)  1,250.  —  Existence 
of  Selenium  Monoxide  (Peirce),  I,  385.  —  Condition  of  Oxidation  of  Man- 
ganese precipitated  by  the  Chlorate  Process  (Gooch  and  Austin),  II,  85.  — 
Action  of  Carbon  Dioxide  on  Soluble  Borates  (Jones),  II,  100.  —  Action  of 
Acetylene  on  Oxides  of  Copper  (Gooch  and  Baldwin),  II,  276.  —  Action  of 
Sodium  Thiosulphate  on  Solutions  of  Metallic  Salts  at  High  Temperatures 
and  Pressures  (Norton),  II,  384. 

ORGANIC  CHEMISTRY. 

Blue  Iodide  of  Starch  (Roberts)  I,  236.  —  Action  of  Urea  and  Sulphocarbanilide 
on  Acid  Anhydrides  (Dunlap),  I,  355.  —  Action  of  Urea  and  Primary  Amines 
on  Maleic  Anhydrides  (Dunlap  and  Phelps),  11,42.  —  Ethers  of  Toluquinone- 
oxime  and  Space  Isomerism  of  Nitrogen  (Bridge  and  Morgan),  II,  145.  — 
Space  Isomerism  of  Toluquinoneoxime  Ethers  (Morgan),  II,  283.  —  Ethers  of 
Isonitrosoguaiacol  and  Space  Isomerism  of  Nitrogen  (Bridge  and  Morgan), 
11,304. 

MINERALOGICAL  CHEMISTRY. 

Rhodochrosite  from  Franklin  Furnace  (Browning),  I,  57.  —  So-called  Perofskite 
from  Magnet  Cove  (Mar),  I,  60. 

ANALYTICAL  CHEMISTRY. 
QUALITATIVE  ANALYSIS. 

Detection  of  Iodine,  Bromine  and  Chlorine  (Gooch  and  Brooks),  I,  47.  —  Detec- 
tion of  Strontium  and  Calcium  (Browning),  1, 121.  —  Detection  of  Arsenic 
with  Antimony  and  Tin  (Gooch  and  Hodge),  I,  231.  —  Detection  of  Per- 
chlorates  (Gooch  and  Kreider),  I,  246.  —  Reduction  of  Arsenic  Acid  (Gooch 
and  Phelps),  I,  265.  —  Separation  and  Identification  of  Potassium  and  Sodium 
(Kreider  and  Breckenridge),  I,  401.  —  Detection  of  Sulphides,  Sulphates, 
Sulphites,  Thiosulphates  (Browning  and  Howe),  II,  134.  —  Separation  of 
Nickel  from  Cobalt  (Browning  and  Hart  well),  II,  344. 


396  SYSTEMATIC  INDEX. 

QUANTITATIVE  ANALYSIS. 

Colorimetric  Methods. 

Detection  and  Approximative  Estimation  of  Minute  Amounts  of  Arsenic  in 
Copper  (Gooch  and  Moseley),  I,  272. 

Electrolytic  Methods. 

Determination  of  Halogens  in  mixed  Silver  Salts  (Gooch  and  Fairbanks),  I,  290. 
Spectroscopic  Methods. 

Determination  of  Potassium  (Gooch  and  Hart),  I,  92.  —  Determination  of  Ru- 
bidium (Gooch  and  Phinney),  I,  157. 

Gravimetric  Methods. 

Determination  of  Chlorine  in  Alkaline  Chlorides  and  Iodides  (Gooch  and  Mar), 

I,  18.  —  Determination  of  Bromine  in  Alkaline  Bromides  and  Iodides  (Gooch 
and  Ensign),  I,  37.  —  Estimation  of  Barium  as  Sulphate  (Mar),  I,  63.  —  Sepa- 
ration of  Strontium  from  Calcium  (Browning),  I,  107.  —  Separation  of  Barium 
from  Calcium  (Browning),  I,  116.  —  Determination  of  Barium  in  presence  of 
Calcium  and  Magnesium  (Mar),  I,  125.  —  Separation  of  Barium  from  Stron- 
tium (Browning),  I,  168.  —  Influence  of  Nitric  Acid  and  Aqua  Regia  on  the 
Precipitation  of  Barium  as   Sulphate  (Browning),  I,   181.  —  Treatment  of 
Barium  Sulphate  (Phiuney),  I,  187.  —  Separation  of  Copper  from  Cadmium 
(Browning),  I,  226. —  Determination  of  Potassium  (Kreider),  I,  282.  —  De- 
termination of  Carbon  Dioxide  (Gooch  and  Phelps),  I,  302.  —  Determination 
of  Selenium  (Peirce),!,  365.  — Estimation  of  Cadmium  as  Oxide  (Browning 
and  Jones),  1, 409.  —  Separation  of  Aluminum  from  Iron  (Gooch  and  Havens), 

II,  20.  —  Separation  of  Aluminum  and  Beryllium  (Havens),  II,  47.  —  Esti- 
mation of  Manganese  as  Sulphate  and  as  Oxides  (Gooch  and  Austin),  II, 
77.  —  Estimation  of  Manganese  separated  as  Carbonate  (Austin),  II,  96.  — 
Separations  of  Aluminum  by  Hydrochloric  Acid  (Havens),  II,  106.  —  De- 
termination of  Manganese  as  Pyrophosphate  (Gooch  and  Austin),  II,  121.  — 
Separation  of  Nickel  and  Cobalt  by  Hydrochloric  Acid  (Havens),  II,  141.  — 
Estimation  of  Boric  Acid  (Gooch  and  Jones),  II,  172.  —  Ammonium  Magne- 
sium Phosphate  of  Analysis  (Gooch  and  Austin),  II,  190.  — Volatilization 
of  Iron  Chlorides  and  Separation  of  Oxides  of  Iron  and  Aluminum  (Gooch 
and    Havens),   II,   215.  —  Double  Ammonium    Phosphates    of    Beryllium, 
Zinc,   Cadmium  (Austin),  II,  252.  —  Separation  of   Iron  from  Chromium, 
Zirconium,  Beryllium,  by  Gaseous  Hydrochloric  Acid  (Havens  and  Way),  II, 
266.  —  Ammonium  Magnesium  Arseniate  of  Analysis  (Austin),  II,  209.  — 
Estimation  of  Thallium  as  Acid   and  Neutral  Sulphates   (Browning),  II, 
317.  —  Separation  and  Determination  of  Mercury  as  Oxalate  (Peters),  II, 
325.  —  Sulphocyanides  of  Copper  and  Silver  (Van  Name),  II,  359.  —  Estima- 
tion of  Caesium  and  Rubidium  as  the  Acid  Sulphates  and  of  Sodium  and 
Potassium  as  the  Pyrosulphates  (Browning),  II,  368. 

Volumetric  Methods. 

Standard  Solutions.  Tartar  Emetic  (Gruener),  I,  216.  Potassium  Permanganate 
(Roberts),  I,  269. 

lodometric  Processes.  Determination  of  Iodine  in  Haloid  Salts  (Gooch  and  Brown- 
ing), 1, 1.  —  Reduction  of  Arsenic  Acid  (Gooch  and  Browning),  I,  30.  —  De- 


SYSTEMATIC  INDEX.  397 

termination  of  Antimony  and  its  Condition  of  Oxidation  (Gooch  and 
Gruener),  1, 73.  —  Estimation  of  Chlorates  (Gooch  and  Smith),  I,  82.  —  Sepa- 
ration of  Antimony  from  Arsenic  (Gooch  and  Banner),  I.  36.  —  Determina- 
tion of  Nitrates  (Gooch  and  Gruener),  I,  132.  —  Determination  of  Iodine  in 
Haloid  Salts  by  action  of  Arsenic  Acid  (Gooch  and  Browning),  I,  173.  — 
Determination  of  Nitrates  (Gruener),  I,  193.  —  Estimation  of  Chlorates  and 
Nitrates,  and  of  Nitrites  and  Nitrates  (Roberts),  1, 219.  —  Estimation  of  Tellu- 
ric Acid  (Gooch  and  Rowland),  I,  277.  —  Reduction  of  Acids  of  Selenium  by 
Hydriodic  Acid  (Gooch  and  Reynolds),  I,  310.  —  Determination  of  Per- 
chlorates  (Kreider),  I,  316.  —  Reduction  of  Selenic  Acid  (Gooch  and  Evans), 

I,  331. —Reduction  of  Selenic  Acid  (Gooch  and  Scoville),  I,  335.  —  Deter- 
mination of  Selenious  and  Selenic  Acids  (Gooch  and  Peirce),  I,  338.  —  Inter- 
action of  Chromic  and  Arsenious  Acids  (Browning),  I,  344.  —  Separation  of 
Selenium  from  Tellurium  (Gooch  and  Peirce),  1, 348.  —  Determination  of  Car- 
bon Dioxide  (Phelps),  I,  369.  —  Estimation  of  Molybdic  Acid  (Gooch  and  Fair- 
banks), I,  375.  — Determination  of  Phosphorus  in  Iron  (Fairbanks),  I,  391.  — 
Reduction  of  Vanadic  Acid  (Browning),  I,  397.  —  Estimation  of  Vanadium 
(Browning  and  Goodman),  II,  4.  —  Determination  of  Oxygen  in  Air  and 
Aqueous  Solution  (Kreider),  II,  11.  — Estimation  of  Molybdenum  (Gooch), 

II,  27.  —  Application  of   lodic  Acid  to  Analysis  of  Iodides   (Gooch  and 
Walker),  II,  33.— Titration  of  Sodium  Thiosnlphate  with  lodic  Acid  (Walker), 
II,   52.  —  Determination  of  Molybdenum  (Gooch  and  Norton),  II,  111.  — 
Analysis  of  Alkalies  and  Acids  (Walker  and  Gillespie),  II,  162.  —  Influence 
of  Hydrochloric  Acid  in  Titrations  by  Thiosulphate,  and  Estimation  of  Se- 
lenious Acid  (Norton),  II,  206.  —  Estimation  of  Iron  in  the  Ferric  State 
(Norton),  II,  230.  —  Determination  of  Tellurous  Acid  in  presence  of  Haloid 
Salts  (Gooch  and  Peters),  II,  238.  —  Estimation  of  Boric  Acid  (Jones),  II, 
244.  — Determination  of  Gold  (Gooch  and  Morley),  II,  269.  —  Estimation  of 
Cerium    (Browning),    II,  289.  —  Estimation    of    Thallium  (Browning  and 
Hutchins),  II,  300.  —  Titration  of  Mercury  by  Sodium  Thiosulphate  (Norton), 
II,  328. —  Estimation  of  Arsenic  Acid  (Gooch  and  Morris),  II,  336. 

Alkalimetric  Processes.    Estimation  of  Boric  Acid  (Jones),  II,  182. 

Oxidimetric  Processes.  Determination  of  Selenious  Acid  (Gooch  and  demons), 
I,  297.  —  Titration  of  Oxalic  Acid  in  presence  of  Hydrochloric  Acid  (Gooch 
and  Peters),  II,  222.  — Determination  of  Tellurous  Acid  in  presence  of  Haloid 
Salts  (Gooch  and  Peters),  II,  238.  —  Separation  and  Determination  of  Mercury 
as  Mercurous  Oxalate  (Peters),  II,  320.  —  Estimation  of  Copper  as  Oxalate, 
with  Separations  (Peters),  II,  347. —  Estimation  of  Calcium,  Strontium,  and 
Barium  as  the  Oxalates  (Peters),  II,  373. 

Precipitation  Processes.  Determination  of  Tellurium  by  Precipitation  as  the 
Iodide  (Gooch  and  Morgan),  II,  1. 

Gasometric  Processes.  Reduction  of  Nitric  Acid  by  Ferrous  Salts  (Roberts), 
I,  203.  —  Estimation  of  Chlorates  and  Nitrates,  and  of  Nitrites  and  Nitrates 
(Roberts),  I,  219. 


INDEX  OF  AUTHORS. 


AUSTIN,  MARTHA.     Estimation  of  Manganese  as  Sulphate  and  Oxide 

(with  Gooch,  F.  A.) H,  77 

Condition  of  Oxidation  of  Manganese  precipitated  by  Chlorate 

Process  (with  Gooch,  F.  A.) II,  85 

Estimation  of  Manganese  Separated  as  Carbonate II,  96 

Determination  of  Manganese  as  Pyrophosphate  (with  Gooch, 

F.  A.) II,  121 

Constitution  of  Ammonium  Magnesium  Phosphate  of  Analysis  (with 

Gooch,  F.  A.) II,  190 

Double  Ammonium  Phosphates  of  Beryllium,  Zinc,  Cadmium,  in 

Analysis II,  252 

Constitution  of  Ammonium  Magnesium  Arseniate  of  Analysis  .  .  II,  309 
BALDWIN,  DEFOREST.  Action  of  Acetylene  on  Oxides  of  Copper  (with 

Gooch,  F.  A.) II,  276 

BRECKENRIDGE,  J.  E.  Separation  and  Identification  of  Potassium  and 

Sodium  (with  Kreider,  D.  Albert) I,  401 

BRIDGE,  JOHN  L.  Ethers  of  Toluquinoneoxime,  and  Space  Isomerism 

of  Nitrogen  (with  Morgan,  Wm.  Conger)  II,  145 

Ethers  of  Isonitrosoguaiacol  and  Space  Isomerism  of  Nitrogen  (with 

Morgan,  Wm.  Conger) II,  304 

BROOKS,  F.  T.  Detection  of  Iodine,  Bromine,  and  Chlorine  (with 

Gooch,  F.  A.) I,  47 

BROWNING,  PHILIP  E.  Determination  of  Iodine  in  Haloid  Salts  (with 

Gooch,  F.  A.)  I,  1 

Reduction  of  Arsenic  Acid  in  Analysis  (with  Gooch,  F.  A.)  ...  I,  30 

Analysis  of  Rhodochrosite  from  Franklin  Furnace I,  57 

Quantitative  Separation  of  Strontium  from  Calcium  by  Amyl 

Alcohol  on  Nitrates I,  107 

Quantitative  Separation  of  Barium  from  Calcium  by  Amyl  Alcohol 

on  Nitrates I,  116 

Separation  and  Detection  of  Strontium  and  Calcium  by  Amyl 

Alcohol  on  Nitrates .  . 1,121 

Quantitative  Separation  of  Barium  from  Strontium  by  Amyl  Alcohol 

on  Bromides I,  168 

Determination  of  Iodine  in  Haloid  Salts  by  Action  of  Arsenic  Acid 

(with  Gooch,  F.  A.) I,  173 

Influence  of  Nitric  Acid  and  Aqua  Regia  on  Precipitation  of  Barium 

as  Sulphate I,  181 

Separation  of  Copper  from  Cadmium  by  Iodide  Method  ....  I,  226 

Interaction  of  Chromic  and  Arsenious  Acids I,  344 

Reduction  of  Vanadic  Acid  by  Hydrobromic  and  Hydriodic  Acids, 

and  Estimation  by  Iodine I,  397 


400  INDEX  OF  AUTHORS. 

VOL.  PAGE 

Estimation  of  Cadmium  as  Oxide  (with  Jones,  Louis  C.)    .    .    .    .      I,  409 
Application  of  Organic  Acids  to  Estimation  of  Vanadium  (with 

Goodman,  Richard  J.) II,     4 

Detection  of  Sulphides,   Sulphates,   Sulphites,  and  Thiosulphates 

(with  Howe,  Ernest) II,  134 

Volumetric  Estimation  of  Cerium  (with  Hanford,  G.  A. ;  Hall,  F.  J. ; 

Cutter,  Wm.  D. ;  Lynch,  Leo  A.) II,  289 

Estimation  of  Thallium  as  Chromate  (with  Hutchius,  George  P.)    .    II,  300 

Estimation  of  Thallium  as  Acid  and  Neutral  Sulphates 11,317 

Qualitative  Separation  of  Nickel  from  Cobalt  by  Ammonia  on  the 

Ferricyanides  (with  Hartwell,  John  B.) II,  344 

On  the  Estimation  of  Caesium  and   Rubidium  as  the  Acid  Sul- 
phates, and  of  Potassium  and  Sodium  as  the  Pyrosulphates     .     .    II,  368 
CLEMONS,  C.  F.    Determination  of  Selenious  Acid  by  Potassium  Perman- 
ganate (with  Gooch,  F.  A.) I,  297 

CUTTER,  WM.  D.    Volumetric  Estimation  of  Cerium  (with  Browning, 

Philip  E.) 11,294 

DANNER,  E.  VV.  Separation  of  Antimony  from  Arsenic  by  Hydrochloric 

and  Hydriodic  Acids  (with  Gooch,  F.  A.) I,    86 

Interaction  of  Potassium  Permanganate  and  Sulphuric  Acid  (with 

Gooch,  F.  A.) I,  145 

DONLAP,  FREDERICK  L.    Action  of  Urea  and  Sulphocarbanilide  on  Acid 

Anhydrides 1, 355 

Action  of  Urea  and  Primary  Amines  on  Maleic  Anhydride  (with 

Phelps,  Isaac  K.) 11,42 

ENSIGN,  J.  R.    Determination  of  Bromine  in  Alkaline  Bromides  and 

Iodides  (with  Gooch,  F.  A.) I,  37 

EVANS,  P.  S.,  JR.     Reduction  of  Selenic  Acid  by  Hydrochloric  Acid 

(with  Gooch,  F.  A.) 1,331 

FAIRBANKS,  CHARLOTTE.    Estimation  of  Halogens  in  Silver  Salts  (with 

Gooch,  F.  A.) .     .      1,290 

lodometric  Estimation  of  Molybdic  Acid  (with  Gooch,  F.  A.)      .     .      I,  375 

lodometric  Determination  of  Phosphorus  in  Iron I,  391 

GILLESPIE,  DAVID  H.  M.    Iodine  in  Analysis  of  Acids  and  Alkalies 

(with  Walker,  Claude  F.) II,  162 

GOOCH,  F.  A.  Determination  of  Iodine  in  Haloid  Salts  (with  Browning, 

P.E.) I,     1 

Determination  of  Chlorine  in  Alkaline  Chlorides  and  Iodides  (with 

Mar,  F.  W.) I,    18 

Reduction  of  Arsenic  Acid  in  Analysis  (with  Browning,  P.  E.)  .     .      I,    30 
Determination  of  Bromine  in  Alkaline  Bromides  and  Iodides  (with 

Ensign,  J.  R.) I,    37 

Detection  of  Iodine,  Bromine,  and  Chlorine  (with  Brooks,  F.  T.)    .      I,    47 
Determination  of  Antimony  and  its  Condition  of  Oxidation  (with 

Gruener,  H.  W.)       I,    73 

Estimation  of  Chlorates  (with  Smith,  C.  G.) I,    82 

Separation  of  Antimony  from  Arsenic  by  Hydrochloric  and  Hydri- 
odic Acids  (with  Danner,  E.  W.) I,    86 

Detection  and  Determination  of  Potassium  Spectroscopically  (with 

Hart,  T.  S.) I,    92 

lodometric  Determination  of  Nitrates  (with  Gruener,  H.  W.)     .    .      I,  132 
Laboratory  Apparatus I,  141 


INDEX  OF  AUTHORS.  401 

VOL.  PAGE 

Interaction  of  Potassium  Permanganate  and  Sulphuric  Acid  (with 

Banner,  E.  W.) I,  145 

Quantitative  Determination  of  Rubidium  by  the  Spectroscope  (with 

Phinney,  J.  I.) I,  157 

Determination  of  Iodine  in  Haloid  Salts  by  Action  of  Arsenic  Acid 

( with  Browning,  P.  E.) 1,173 

Detection  and  Separation  of  Arsenic  with  Antimony  and  Tin  (with 

Hodge,  B.) 1,231 

Detection  of  Alkaline  Perchlorates  with  Chlorides,  Chlorates,  and 

Nitrates  (with  Kreider,  D.  Albert)  I,  246 

Generation  of  Chlorine  (with  Kreider,  D.  Albert) I,  260 

Reduction  of  Arsenic  Acid  by  Hydrochloric  Acid  and  Potassium 

Bromide  (with  Phelps,  I.  K.) I,  265 

Detection  and  Estimation  of  Minute  Amounts  of  Arsenic  in  Copper 

(with  Moseley,  H.  P.) I,  272 

lodometric  Estimation  of  Telluric  Acid  (with  Rowland,  J.)  ...  I,  277 
Estimation  of  Halogens  in  Silver  Salts  (with  Fairbanks,  Charlotte)  I,  290 
Determination  of  Selenious  Acid  by  Potassium  Permanganate  (with 

demons,  C.  F.) I,  297 

Precipitation  and  Gravimetric  Determination  of  Carbon  Dioxide 

(with  Phelps,  I.  K.) I,  302 

Reduction  of  Acids  of  Selenium  by  Hydriodic  Acid  (with  Reynolds, 

W.G.) 1,310 

Reduction  of  Selenic  Acid  by  Hydrochloric  Acid  (with  Evans,  P.  S., 

Jr.) 1,331 

Reduction  of  Selenic  Acid  by  Potassium  Bromide  in  Acid  Solution 

(witb  Scoville,  W.  S.) I,  335 

lodometric  Determination  of  Selenious  and  Selenic  Acids  (with 

Peirce,  A.  W.) I,  338 

Separation  of  Selenium  from  Tellurium  by  difference  in  Volatility 

of  Bromides  (with  Peirce,  A.  W.) I,  348 

lodometric  Estimation  of  Molybdic  Acid  (with  Fairbanks,  Char- 
lotte)   1,375 

Determination  of  Tellurium  by  precipitation  as  Iodide  (with  Mor- 
gan, W.  C.) II,  1 

Separation  of  Aluminum  from  Iron  (with  Havens,  F.  S.)  .  .  .  .  II,  20 

Estimation  of  Molybdenum  lodometrically II,  27 

Application  of  lodic  Acid  to  Analysis  of  Iodides  (with  Walker, 

C.  F.) II,  33 

Estimation  of  Manganese  as  Sulphate  and  Oxide  (with  Martha 

Austin)  II,  77 

Condition  of  Oxidation  of  Manganese  precipitated  by  Chlorate 

Process  (with  Austin,  Martha) II,  85 

lodometric  Determination  of  Molybdenum  (with  Norton,  John  T., 

Jr.) 11,111 

Determination  of  Manganese  as  Pyrophosphate  (with  Austin, 

Martha) II,  121 

Estimation  of  Boric  Acid  (with  Jones,  Louis  Cleveland)  .  .  .  .  II,  172 
Constitution  of  Ammonium  Magnesium  Phosphate  of  Analysis 

(with  Austin,  Martha)  II,  190 

Volatilization  of  Iron  Chlorides,  and  Separation  of  Oxides  of  Iron 

and  Aluminum  (with  Havens,  Franke  Stuart) II,  215 

VOL.  n.  —  26 


402  INDEX  OF  AUTHORS. 

VOL.  PAGE 

Titration  of  Oxalic  Acid  by  Potassium  Permanganate  in  presence 

of  Hydrochloric  Acid  (with  Peters,  C.  A.) II,  222 

Determination  of  Tellurous  Acid  in  presence  of  Haloid  Salts  (with 

Peters,  C.  A.) II,  238 

lodometric  Determination  of  Gold  (with  Morley,  Frederick  H.)       .    II,  269 
Action  of  Acetylene  on  Oxides  of  Copper  (with  Baldwin,  De  Forest)     II,  276 
lodometric  Estimation  of  Arsenic  Acid  (with  Morris,  Julia  C.)  .    .    II,  236 
GOODMAN,  RICHARD  J.    Application  of  Organic  Acids  to  Estimation  of 

Vanadium  (with  Browning,  Philip  E.) II,      4 

GRUBNER,  H.  W.    Determination  of  Antimony  and  its  condition  of  Oxi- 
dation (with  Gooch,  F.  A.) I,    73 

lodometric  Determination  of  Nitrates  (with  Gooch,  F.  A.)       ...      I,  132 

lodometric  Determination  of  Nitrates I,  193 

Stability  of  Standard  Solutions  of  Tartar  Emetic I,  216 

HALL,  F.  J.      Volumetric    Estimation   of    Cerium  (with  Browning, 

Philip  E.)       II,  290 

HANFORD,  G.  A.    Volumetric  Estimation  of  Cerium  (with  Browning, 

Philip  E.) 11,290 

HART,  T.  S.    Detection  and  Determination  of  Potassium  Spectroscopi- 

cally  (with  Gooch,  F.  A.) I,    92 

HARTWELL,  JOHN  B.    Qualitative  Separation  of  Nickel  from  Cobalt  by 

Ammonia  on  the  Ferricyanides  (with  Browning,  Philip  E.)     .     .     II,  344 
HAVENS,  FRANKE  STUART.    Separation  of  Aluminum  from  Iron  (with 

Gooch,  F.  A.) II,    20 

Separation  of  Aluminum  and  Beryllium  by  Hydrochloric  Acid  .     .    II,    47 
Further  Separations  of  Aluminum  by  Hydrochloric  Acid    .    .    .     .    II,  106 

Separation  of  Nickel  and  Cobalt  by  Hydrochloric  Acid II,  141 

Volatilization  of  Iron  Chlorides,  and  Separation  of  Oxides  of  Iron 

and  Aluminum  (with  Gooch,  F.  A.)       II,  215 

Separation  of  Iron  from  Chromium,  Zirconium,  and  Beryllium  by 
Action  of  Gaseous  Hydrochloric  Acid  on  the  Oxides  (with  Way, 

Arthur  Fitch) II,  266 

HODGE,  B.    Detection  and  Separation  of  Arsenic  with  Antimony  and 

Tin  (with  Gooch,  F.  A.)        I,  231 

HOWE,  ERNEST.    Detection  of   Sulphides,   Sulphates,  Sulphites,  and 

Thiosulphates  (with  Browning,  Philip  E.) II,  134 

HOWLAND,  J.    lodometric  Estimation  of  Telluric  Acid  (with  Gooch, 

F.  A.) I,  277 

HUTCHINS,  GEORGE  P.    Estimation  of  Thallium  as  Chromate  (with 

Browning,  Philip  E.) II,  300 

JONES,  Louis  CLEVELAND.    Estimation  of  Cadmium  as  Oxide  (with 

Browning,  Philip  E.) I,  409 

Action  of  Carbon  Dioxide  on  Soluble  Borates II,  100 

Estimation  of  Boric  Acid  (with  Gooch,  F.  A.) II,  172 

Volumetric  Estimation  of  Boric  Acid 11,182 

lodometric  Estimation  of  Boric  Acid .    II,  244 

KREIDER,  D.  ALBERT.    Detection  of  Alkaline  Perchlorates  with  Chlo- 
rides, Chlorates,  and  Nitrates  (with  Gooch,  F.  A.) I,  246 

Generation  of  Chlorine  (with  Gooch,  F.  A.)        I,  260 

Preparation  of  Perchloric  Acid  and  Determination  of  Potassium     .      I,  282 

Laboratory  Apparatus I,  306 

Quantitative  Determination  of  Perchlorates I,  316 


INDEX  OF  AUTHORS.  403 

VOL.  PAGE 

Separation  and  Identification  of    Potassium    and  Sodium   (with 

Breckenridge,  J.  E.) «.    .      I,  401 

Determination  of  Oxygen  in  Air  and  Aqueous  Solution      ....    II,    1 1 
LYNCH,  LEO  A.    Volumetric  Estimation  of  Cerium  (with  Browning, 

Philip  E.)       II,  297 

MAR,  F.  W.     Determination  of  Chlorine  in  Alkaline  Chlorides  and 

Iodides  (with  Gooch,  F.  A.) I,    18 

So-called  Perofskite  from  Magnet  Cove I,    60 

Estimation  of  Barium  as  the  Sulphate I,    63 

Determination  of  Barium  in  presence  of  Calcium  and    Magne- 
sium   I,  125 

MORGAN,  WM.  CONGER.    Determination  of  Tellurium  hy  precipitation 

as  Iodide  (with  Gooch,  F.  A.) II,      1 

Ethers  of  Toluquinoneoxime  and  Space  Isomerism  of  Nitrogen 

(with  Bridge,  John  L.) II,  145 

Space  Isomerisms  of  Toluquinoneoxime  Ethers II,  283 

Ethers  of   Isonitrosoguaiacol  and  Space  Isomerism  of  Nitrogen 

(with  Bridge,  John  L.)        II,  304 

MORLEY,  FREDERICK  H.    lodometric  Determination  of   Gold   (with 

Gooch,  F.  A.) H,  269 

MORRIS,   JULIA    C.    lodometric    Estimation  of   Arsenic  Acid   (with 

Gooch,  F.  A.) II,  336 

MOSELEY,   H.  P.    Detection  and  Estimation  of  Minute  Amounts  of 

Arsenic  in  Copper  (with  Gooch,  F.  A.) I,  272 

NORTON,  JOHN  T.,  Jr.      lodometric  Determination  of  Molybdenum 

(with  Gooch,  F.  A.) II,  111 

Hydrochloric  Acid  in  Titrations  by  Sodium  Thiosulphate,  and  Esti- 
mation of  Selenious  Acid II,  206 

Estimation  of  Iron  in  Ferric  Condition  by  Sodium  Thiosulphate 

and  Iodine II,  230 

Titration  of  Mercury  by  Sodium  Thiosulphate II,  328 

The  Action  of  Sodium  Thiosulphate  on  Solutions  of  Metallic  Salts 

at  High  Temperatures  and  Pressures II,  384 

PEIRCE,  A.  W.    lodometric  Determination  of  Selenious  and  Selenic 

Acids  (with  Gooch,  F.  A.) I,  338 

Separation  of  Selenium  from  Tellurium  by  difference  in  Volatility 

of  Bromides  (with  Gooch,  F.  A.)       I,  348 

Gravimetric  Determination  of  Selenium I,  365 

Existence  of  Selenium  Monoxide I,  385 

PETERS,  CHARLES  A.  Titration  of  Oxalic  Acid  by  Potassium  Permangan- 
ate in  presence  of  Hydrochloric  Acid  (with  Gooch,  F.  A.)     ...    II,  222 
Determination  of  Tellurous  Acid  in  presence  of  Haloid  Salts  (with 

Gooch,  F.  A.) 11,238 

Determination  of  Mercury  as  Mercurous  Oxalate H,  320 

Volumetric  Estimation  of  Copper  with  Separation  from  Cadmium, 

Arsenic,  Tin,  Iron,  and  Zinc II,  347 

The  Estimation  of  Calcium,  Strontium,  and  Barium  as  the  Oxalates.    II,  373 
PHELPS,  ISAAC  K.   Reduction  of  Arsenic  Acid  by  Hydrochloric  Acid 

Potassium  Bromide  (with  Gooch,  F.  A.) I,  265 

Precipitation  and  Gravimetric  Determination  of  Carbon  Dioxide 

(with  Gooch,  F.  A.) I,  302 

lodometric  Determination  of  Carbon  Dioxide I,  369 


404  INDEX  OF  AUTHORS. 

VOL.  PAGE 

Action  of  Urea  and  Primary  Amines  on  Maleic  Anhydride  (with 

Duulap,  Frederick  L.) II,  42 

Combustion  of  Organic  Substances  in  the  Wet  Way II,  62 

PHINNEY,  J.  L  Quantitative  Determination  of  Rubidium  by  the  Spec- 
troscope (with  Gooch,  F.  A.) I,  157 

Treatment  of  Barium  Sulphate  in  Analysis I,  187 

REYNOLDS,  W.  G.    Reduction  of  Acids  of  Selenium  by  Hydriodic  Acid 

( with  Gooch,  F.  A.) 1,310 

ROBERTS,  CHARLOTTE  F.    Reduction  of  Nitric  Acid  by  Ferrous  Salts      I,  203 
Estimation  of  Chlorates  and  Nitrates,  and  of  Nitrites  and  Nitrates  .      I,  219 

Blue  Iodide  of  Starch I,  236 

Action  of  Reducing  Agents  on  lodic  Acid I,  250 

Standardization  of  Potassium  Permanganate  in  Iron  Analysis   .    .      I,  269 
SCOVILLE,  W.  S.    Reduction  of  Selenic  Acid  by  Potassium  Bromide  in 

Acid  Solution  (with  Gooch,  F.  A.) I,  335 

SMITH,  C.  G.    Estimation  of  Chlorates  (with  Gooch,  F.  A.)    ....      I,    82 
VAN  NAME,  R.  G.    The  Sulphocyanides  of  Copper  and  Silver  in  Gravi- 
metric Analysis       II,  359 

WALKER,  CLAUDE.    Application  of  lodic  Acid  to  Analysis  of  Iodides 

(with  Gooch,  F.  A.) II,    33 

Titration  of  Sodium  Thiosulphate  with  lodic  Acid II,    52 

Iodine  in  Analysis  of  Acids  and  Alkalies  (with  Gillespie,  David 

H.  M.) II,  162 

WAY,  ARTHUR  FITCH.  Separation  of  Iron  from  Chromium,  Zir- 
conium, and  Beryllium,  by  action  of  Gaseous  Hydrochloric  Acid 
on  the  Oxides  (with  Havens,  Franke  Stuart) II,  266 


INDEX  OF  SUBJECTS. 

VOL.  PAGE 

Acetylene,  action  of,  on  oxides  of  copper  (Gooch  and  Baldwin)    .    .    .  11,276 

Acids,  application  of  iodine  in  analysis  of  (Walker  and  Gillespie)    .     .  II,  162 

of  selenium,  reduction  of,  by  hydriodic  acid  (Gooch  and  Reynolds)  .  I,  310 

Acid  anhydrides,  action  of  urea  and  sulphocarbanilide  upon  (Dunlap)    .  I,  355 

Alkalies,  application  of  iodine  to  analysis  of  (Walker  and  Gillespie)    .  II,  162 

Alkaline  bromides,  determination  of  bromine  in  (Gooch  and  Ensign)    .  I,    37 
Alkaline  chlorides,  determination  of  chlorine  in  alkaline  iodides  mixed 

with  (Gooch  and  Mar) I,    18 

Alkaline  iodides,   determination  of  chlorine  in  (Gooch  and  Mar)   .     .  I,    18 

determination  of  bromine  in  (Gooch  and  Ensign)    ......  I,    37 

Alkaline  perchlorates,  detection  of,  associated  with  chlorides,  chlorates, 

and  nitrates  (Gooch  and  Kreider) • 1,246 

Aluminum  salts,  action  of  sodium  thiosulphate  upon,  at  high  tempera- 
tures and  pressures  (Norton) II,  388 

Aluminum,  separation  of,  by  hydrochloric  acid,  from  iron  (Gooch  and 

Havens) II,    20 

separation  of,  by  hydrochloric  acid,  from  beryllium  (Havens)    .    .  II,    47 

separation  of,  by  hydrochloric  acid,  from  bismuth,  copper,  and  .     .  . 

mercury  (Havens) II,  109 

separation  of,  by  hydrochloric  acid,  from  zinc  (Havens)    ....  II,  107 

Aluminum  oxide,  separation  of  oxides  of  iron  from  (Gooch  and  Havens)  II,  215 

Ammonium  magnesium  arseniate  in  analysis,  constitution  of  (Austin)    .  II,  309 

Ammonium  beryllium  phosphate  in  analysis  (Austin) II,  253 

Ammonium  cadmium  phosphate  in  analysis  (Austin) II,  262 

Ammonium  magnesium  phosphate  in  analysis  (Gooch  and  Austin)    .     .  11,190 

Ammonium  zinc  phosphate  in  analysis  (Austin) II,  257 

Amyl  alcohol,  use  of,  in  detecting  strontium  and  calcium  (Browning)  .  I,  121 


use  of,  in  separating  strontium  and  calcium  (Browning) 

use  of,  in  separating  barium  and  calcium  (Browning) 

use  of,  in  separating  barium  and  strontium  (Browning)       .... 

Antimony,  detection  of  arsenic  associated  with  (Gooch  and  Hodge)    .    . 

determination  of,  and  its  condition  of  oxidation  (Gooch  and  Gruener) 

separation  of,  from   arsenic,   by  hydrochloric  and  hydriodic  acids 


(Gooch  and  Banner) 


Antimonious  chloride,  decomposition  of  nitrates  by  (Gruener)  .... 
Antimouic  acid,  salts  of,  reduced  by  potassium  iodide  and  sulphuric  acid, 


and  estimated  iodometrically  (Gooch  and  Gruener) 


Apparatus  —  burette  clip  (Gooch) 


121 

116 

^168 

231 

73 


86 
199 

73 
141 

,264 
,308 


chlorine  generator  (Gooch  and  Kreider) 

force  pump  (Kreider) 

hot  filter  (Kreider) I,  306 

mercury  washer  (Gooch) 1,143 

steam  evaporator  (Gooch)        I,  142 

support  (Gooch) I,  142 


406  INDEX  OF  SUBJECTS. 

VOL.  PAGE 

valve  (Kreider)       1, 307 

used  in  analysis  of  iodides  by  iodic  acid  (Gooch  and  Walker)  ...  II,    37 

used  in  combustion  of  organic  substances  in  the  wet  way  (Phelps    .  II,    68 
used  in  estimation  of  carbon  dioxide  gravimetrically  (Gooch  and 

Phelps) 1,302 

used  in  estimation  of  iodine  in  haloid  salts  (Gooch  and  Browning)  .  I,    12 

used  in  estimation  of  molybdenum  (Gooch  and  Norton)      .    .    .    .  11,114 
used  in  estimation  of  molybdic  acid  (Gooch  and  Fairbanks)        .    I,  378,  382 

used  in  estimation  of  molybdic  acid  (Fairbanks) I,  394 

used  in  estimation  of  nitrates  (Gooch  and  Gruener) I,  137 

used  in  estimation  of  oxygen  in  air  and  aqueous  solution  (Kreider)  II,    17 
used  in  estimation  of  selenium  iodometrically,  by  volatilization  of 

the  bromide  (Gooch  and  Peirce) I,  350 

used  in  reduction  of  arsenic  acid  (Gooch  and  Browning)     ....  I,    33 

used  in  reduction  of  nitric  acid  by  ferrous  salts  (Roberts)       .     .    .  1,208 
Arsenic,  detection  of,  associated  with  antimony  and    tin  (Gooch  and 

Hodge)       1,231 

detection  and  approximative  estimation  of,  in  copper  (Gooch  and 

Moseley) 1, 272 

separation  of  antimony  from,  by  hydrochloric  and  hydriodic  acids 

(Gooch  and  Danner) I,    86 

separation  of  copper  as  oxalate  from  (Peters)         II,  347 

Arsenious  acid,  action  of,  upon  cerium  dioxide  (Browning  and  Cutter)  II,  294 

interaction  of,  with  chromic  acid  (Browning)       I,  344 

Arsenic  acid,  determination  of,  by  reduction  with  potassium  iodide  and 
sulphuric  acid,  and  titration  by  iodine  in  alkaline  solution  (Gooch 

and  Browning)         I,    30 

iodometric  estimation  of  (Gooch  and  Morris)    . II,  336 

reduction  of,  in  analysis  (Gooch  and  Browning) II,    30 

reduction  of,  by  action  of  hydrochloric  acid  and  potassium  bro- 
mide (Gooch  and  Phelps) .  1, 265 

use  of,  to  liberate  iodine  in  quantitative  estimation  of  iodides  (Gooch 

and  Browning) I,      1 

use  of,  in  determination  of  iodine  in  haloid  salts  (Gooch  and  Brown- 
ing)       I,      1 

Aqua  regia,  influence  of,  on  the  precipitation  of  barium  as  the  sul- 
phate (Browning)         I,  181 

Barium,  determination  of,  in  presence  of  calcium  and  magnesium  (Mar)  I,  125 

estimation  of,  as  oxalate  (Peters) II,  373 

points  in  estimation  of,  as  sulphate  (Mar)        ....            ...  I,    63 

precipitation  of,  as  sulphate,  in  presence  of  nitric  acid  and  aqua 

regia  (Browning) I,  181 

quantitative  separation  of,  from  calcium  by  amyl  alcohol  on  the  ni- 
trates (Browning) I,  116 

quantitative  separation  of,  from  strontium  by  amyl  alcohol  on  the 

bromides  (Browning) I,  168 

Barium  chlorides,  precipitation  and  separation  of,  from  calcium  and  mag- 
nesium, by  hydrochloric  acid  and  ether  (Mar) I,  125 

Barium  sulphate,  influence  of  hydrochloric  acid  upon  precipitation  of 

(Mar)       I,    63 

purification  of,  by  crystallizing  from  sulphuric  acid  (Mar)    ....  I,    71 

treatment  of,  in  analysis  (Phinney) I,  187 


INDEX  OF  SUBJECTS.  407 

VOL.  PAGE 

Beryllium,  separation  of,  from  aluminum,  by  action  of  hydrochloric  acid 

(Havens) II,    47 

separation  of  iron  from,  by  action  of  hydrochloric  acid  (Havens  and 

Way) 11,266 

Beryllium  ammonium  phosphate  in  analysis  (Austin) II,  253 

Beryllium  salt,  action  of  sodium  thiosulphate  upon,  at  high  tempera- 
tures and  pressures  (Norton) n,  391 

Bismuth,  separation  of  aluminum  from  (Havens) II,  109 

Blue  iodide  of  starch  (Eoberts)       I,  236 

Borates  (soluble),  action  of  carbon  dioxide  on  (Jones)       .  II,  100 

Boric  acid,  estimation  of  (Gooch  and  Jones) II,  172 

iodometric  method  for  estimation  of  (Jones) II}  244 

use  of  calcium  oxide  as  a  retainer  for  (Gooch  and  Jones)    ....  11,175 

use  of  sodium  tungstate  as  a  retainer  for  (Gooch  and  Jones)  .    .    .  II,  178 

volumetric  estimation  of  (Jones) II,  182 

Bromine,  detection  of,  in  presence  of  chlorine  and  iodine  (Gooch  and 

Brooks) I,    47 

,  determination  of,  in  alkaline  bromides  and  iodides  (Gooch  and 

Ensign)       I,    37 

volatilization  of,  from  aqueous  solutions  of  bromide  and  chloride 

by  action  of  sulphuric  acid  and  nitrous  acid  (Gooch  and  Ensign)  .  I,    43 

Cadmium,  estimation  of,  as  oxide  (Browning  and  Jones) I,  409 

Separation  of  copper  from,  by  the  iodide  method  (Browning)      .     .  I,  226 

Separation  of  copper  from,  as  oxalate  (Peters)       II,  354 

Cadmium  ammonium  phosphate,  in  analysis  (Austin) II,  262 

Caesium,  estimation  of,  as  the  acid  sulphate  (Browning) II,  368 

Calcium,  determination  of  barium  in  presence  of  (Mar) I,  125 

estimation  of,  as  oxalate  (Peters) II,  373 

quantitative  separation  of  barium  from,  by  action  of  amyl  alcohol  on 

the  nitrates  (Browning) ' I,  1 1 6 

separation  of,  from  strontium,  and  detection  of,  by  action  of  amyl 

alcohol  on  the  nitrates  (Browning) 1,121 

Calcium  oxide,  use  of,  as  a  retainer  for  boric  acid  (Gooch  and  Jones)      .  II,  175 

Carbon  dioxide,  action  of  on  soluble  borates  (Jones) II,  100 

iodometric  method  for  determination  of  (Phelps) 1,369 

precipitation  and  gravimetric  determination  of  (Gooch  and  Phelps)  I,  302 
Cerium,  modified  Bunsen  method  for  determination  of  (Browning,  Han- 
ford,  and  Hall) II,  290 

volumetric  estimation  of  (Browning) II,  289 

Cerium  dioxide,  action  of  arsenious  acid  upon  (Browning  and  Cutter)     .  II,  294 
Cerium  oxalate,  estimation  of,  by  potassium  permanganate  (Browning 

and  Lynch) 11,297 

Chlorates,  detection  of  -perchlorates  associated  with  (Gooch  and  Kreider)  I,  246 

estimation  of  (Gooch  and  Smith) I,    82 

Chlorates  and  nitrates,  estimation  of,  in  one  operation  (Roberts)    ...  1,219 
Chlorate  process,  condition  of  oxidation  of  manganese  precipitated  in 

(Gooch  and  Austin) II,  85 

Chlorides,  detection  of  perchlorates  associated  with  (Gooch  and  Kreider)  I,  246 
Chlorine,  detection  of,  in  presence  of  bromides  and  iodides  (Gooch  and 

Brooks) I,   47 

direct  determination  of,  in  alkaline  chlorides  and  iodides  (Gooch  and 

Mar) I,    18 


408  INDEX  OF  SUBJECTS. 


generation  of,  by  hydrochloric  acid  and  potassium  chlorate  (Gooch 

and  Kreider) I,  260 

Chromium  salt,  action  of  sodium  thiosulphate  upon,  at  high  tempera- 
tures and  pressures  (Norton) II,  389 

Chromium,  separation  of  iron  from,  by  gaseous  hydrochloric  acid  (Ha- 
vens and  Way) II,  266 

Chromic  acid,  interaction  of,  with  arsenious  acid  (Browning)    ....  I,  344 

use  of,  in  combustion  of  organic  substances  in  the  wet  way  (Phelps)  II,    67 

Cobalt,  separation  of,  from  nickel  (Havens) II,  141 

separation  of  nickel  from,  by  action  of  ammonium  hydroxide  on  the 

ferricyanides  (Browning  and  Hart  well) II,  344 

Combustion  of  organic  substances  in  the  wet  way  (Phelps) II,    62 

Copper,  detection  and  approximate  estimation  of  minute  amounts  of 

arsenic  in  (Gooch  and  Moseley) I,  272 

estimation  of,  as  oxalate,  with  separation  from  cadmium,  arsenic, 

tin,  iron,  and  zinc  (Peters) II,  347 

preparation  of,  free  from  arsenic  (Gooch  and  Moseley) I,  275 

separation  of  aluminum  from  (Havens) II,  109 

separation  of,  from  cadmium  by  the  iodide  method  (Browning)   .     .  I,  226 

Copper  oxides,  action  of  acetylene  on  (Gooch  and  Baldwin) 11,276 

Copper  sulphocyanide  in  gravimetric  analysis  (Van  Name) 11,359 

Dibrommaleinamide,  preparation  of,  from  urea  and  dibrommaleic  anhy- 
dride (Dunlap) I,  358 

Dibrommale'inuric  acid,  preparation  of,  from  urea  and  dibrommaleic  an- 
hydride (Dunlap) 1, 358 

Dibromtoluquinonemetaoxime  beuzoyl  ether  (Bridge  and  Morgan)    .     .  II,  157 
Dibromtoluquinoneorthooxime  benzoyl  ether  (Bridge  and  Morgan)    .     .  II,  161 
Dibromtoluquinoneorthooxime  methyl  ether  (Bridge  and  Morgan)     .     .  II,  159 
Dichlormaleinimide  preparation  of,  from  urea  and  dichlormale'ic  anhy- 
dride (Dunlap) 1, 357 

Dichlormale'inuric  acid,  preparation  of,  by  action  of  urea  on  dichlorma- 

leic  anhydride  (Dunlap) I,  356 

Double  ammonium  phosphates  of  beryllium,  zinc,  and  cadmium  in  ana- 
lysis (Austin) II,  252 

Electrolytic  iron,  use  of,  in  standardizing  permanganate  solutions  (Rob- 
erts)    1,269 

Ethers  of  toluquinoneoxime,  and  their  bearing  on  the  space  isomerism 

of  nitrogen  (Bridge  and  Morgan) II,  145 

Ferric  alum,  use  of,  with  nitric  acid,  to  liberate  iodine  from  haloid  salts 

(Gooch  and  Mar) I,  23 

Ferrous  salts,  use  of,  in  reduction  of  nitric  acid  (Roberts) I,  203 

Gold,  iodometric  determination  of  (Gooch  and  Morley) II,  269 

Halogens,  estimation  of,  in  mixed  silver  salts  (Gooch  and  Fairbanks)     .  I,  290 
Haloid  salts,  determination  of  iodine  in  (Gooch  and  Browning)       ...  I,      2 
determination  of  tellurous  acid  in  presence  of  (Gooch  and  Peters)    .  II,  238 
Hydriodic  acid,  action  of,  with  hydrochloric  acid  in  separation  of  anti- 
mony from  arsenic  (Gooch  and  Danner) I,    86 

use  of,  in  reduction  of  acids  of  selenium  (Gooch  and  Reynolds)   .     .  I,  310 

use  of,  in  reduction  of  vanadic  acid  ( Browning) I,  397 

Hydrobromic  acid,  use  of,  in  reduction  of  vanadic  acid  (Browning)     .     .  I,  397 
Hydrochloric  acid,  action  of,  with  hydriodic  acid,  in  separation  of  anti- 
mony from  arsenic  (Gooch  and  Danner) I,   86 


INDEX  OF  SUBJECTS.  409 

VOL.  PAGE 

influence  of,  upon  the  precipitation  of  barium  sulphate  (Mar)  ...  II,   63 

influence  of,  in  titrations  by  sodium  thiosulphate,  with  special  refer- 
ence to  the  estimation  of  selenious  acid  (Norton) II,  206 

use  of,  in  reduction  of  seleuic  acid  {Gooch  and  Evans) I,  331 

use  of,  in  separation  of  aluminum  from  iron  (Gooch  and  Havens)     .  II,  20 

use  of,  in  separation  of  aluminum  from  zinc,  copper,  mercury,  bis- 
muth (Havens) II,  106 

use  of,  with  ether,  to  precipitate  barium  chloride  in  presence  of  salts 

of  magnesium  and  calcium  (Mar) 1,125 

use  of,  with  potassium  bromide,  in  reducing  and  volatilizing  arsenic 

acid  (Gooch  and  Phelps) I,  265 

use  of,  with  potassium  bromide,  in  separating  arsenic  from  copper 

(Gooch  and  Moseley) I,  272 

use  of,  with  potassium  chlorate,  to  generate  chlorine  (Gooch  and 

Kreider) 1, 260 

use  of,  with  potassium  iodide  in  volatilizing  arsenic  (Gooch  and 

Hodge) 1, 231 

titration  of  oxalic  acid  by  potassium  permanganate  in  presence  of 

(Gooch  and  Peters) 11,222 

volatility  of,  in  aqueous  solutions  containing  sulphuric  acid  and  so- 
dium chloride  (Gooch  and  Mar)      I,    19 

Hydrochloric  acid  (gaseous)  use  of,  in  separation  of  iron  from  chromium, 

zirconium,  and  beryllium  (Havens  and  Way) II,  266 

lodic  acid,  action  of  iodine  on,  in  presence  of  hydrochloric  acid  (Roberts)  I,  257 

action  of  reducing  agents  on,  in  presence  of  hydrochloric  acid  (Rob- 

erts) I,  252 

application  to  the  analysis  of  iodides  (Gooch  and  Walker)  ....  II,  33 

use  of,  in  absorption  of  nitric  oxide  (Roberts) 1,250 

use  of,  in  titration  of  sodium  thiosulphate  (Walker) II,   52 

Iodides,  application  of  iodic  acid  to  the  analysis  of  (Gooch  and  Walker)  .  II,  33 

Iodide  method,  use  of,  in  separating  copper  from  cadmium  (Browning)  .  I,  226 

Iodine,  action  of,  on  iodic  acid  in  presence  of  hydrochloric  acid  (Roberts)  I,  257 

application  of,  in  analysis  of  alkalies  and  acids  (Walker  and  Gilles- 

pie) 11,162 

detection  of,  in  presence  of  chlorine  and  bromine  (Gooch  and 

Brooks) I,    47 

.determination  of,  in  haloid  salts,  by  action  of  arsenic  acid  (Gooch 

and  Browning) I,    1 

liberation  of,  from  haloid  salts,  by  arsenic  acid  (Gooch  and  Brown- 
ing)        I,    1 

liberation  of,  from  haloid  salts,  by  ferric  alum  with  nitric  acid  (Gooch 

and  Mar) I,   23 

liberation  of,  from  haloid  salts,  by  nitrous  acid  (Gooch  and  Mar)      .  I,   27 

use  of,  in  estimating  iron  reduced  from  the  ferric  state  by  sodium 

thiosulphate  (Norton)       II,  230 

lodometric  determination  of  gold  (Gooch  and  Morley) II,  269 

of  molybdenum  (Gooch  and  Norton) 11,111 

of  nitrates  (Gooch  and  Gruener) I,  132 

of  nitrates  (Gruener) 1,193 

of  selenious  and  selenic  acids  (Gooch  and  Peirce) I,  338 

lodometric  estimation  of  alkalies  and  acids  (Walker  and  Gillespie)    .    .  II,  162 

of  antimonic  acid  (Gooch  and  Gruener) I,  73 


410  INDEX  OF  SUBJECTS. 

VOL.  PAGE 

of  antimony  separated  from  arsenic  (Gooch  and  Danner)    ....  I,  86 

of  arsenic  acid  (Gooch  and  Browning) I,  30 

of  arsenic  acid  (Gooch  and  Morris)        II,  336 

of  boric  acid  (Jones) n,  244 

of  carbon  dioxide  (Phelps)       I?  359 

of  chlorates  (Gooch  and  Smith) I,   82 

of  chromic  acid  (Browning) I?  344 

of  cerium  (Browning,  Hanford,  and  Hall) II,  290 

of  gold  (Gooch  and  Morley) II,  269 

of  iodides  (Gooch  and  Walker) II,  33 

of  iodine  in  haloid  salts  (Gooch  and  Browning) I,      i 

of  iron  (Norton) II,  230 

of  mercury  (Norton) II,  328 

of  molybdenum  (Gooch) II,  27 

of  molybdenum  (Gooch  and  Norton) 11,111 

of  molybdic  acid  (Gooch  and  Fairbanks) I,  375 

of  nitrates  ( Gooch  and  Gruener) 1,132 

of  nitrates  (Gruener) I,  193 

of  oxygen,  in  air  and  in  aqueous  solution  (Kreider) II,  11 

of  oxygen,  in  perchlorates  (Kreider) I,  316 

of  phosphorus  in  iron  (Fairbanks) I,  391 

of  selenious  acid  (Gooch  and  Reynolds) I,  310 

of  selenious  acid  (Gooch  and  Peirce) 1,338 

of  selenic  acid  (Gooch  and  Reynolds) 1,314 

of  selenic  acid  (Gooch  and  Peirce) I,  338 

of  selenium  associated  with  tellurium  (Gooch  and  Peirce)  ....  I,  348 

of  tellurous  acid  (Gooch  and  Peters) II,  238 

of  vanadic  acid  (Browning) 1,397 

of  vauadic  acid  (Browning  and  Goodman) II,      4 

lodometric  method  for  the  determination  of  carbon  dioxide  (Phelps)      .  I,  369 

for  the  determination  of  phosphorus  in  iron  (Fairbanks)     ....  I,  391 

for  the  estimation  of  boric  acid  (Jones) II,  244 

Iron,  estimation  of,  in  the  ferric  state  by  reduction  with  sodium  thiosul- 

phate  and  titration  with  iodine  (Norton) II,  230 

iodometric  method  for  determination  of  phosphorus  in  (Fairbanks)  I,  391 
method  for  the  separation  of  aluminum  from  (Gooch  and  Havens)  .  II,    20 
separation  of,  from  chromium,  zirconium,  and  beryllium,  by  gase- 
ous hydrochloric  acid  ( Havens  and  Way) II,  266 

separation  of  copper  oxalate  from  (Peters) II,  347 

Iron  analysis,  standardization  of  potassium  permanganate  in  (Roberts) .  I,  269 

Iron  oxides,  separation  of,  from  aluminum  oxide  (Gooch  and  Havens)    .  II,  215 

Isonitrosoguaiacol,  and  salts  of  (Bridge  and  Morgan) II,  306 

Isonitrosoguaiacol  benzoyl  ether  (Bridge  and  Morgan) II,  307 

Isonitrosoguaiacol  benzoyl  ether  dibromide  (Bridge  and  Morgan)       .     .  II,  307 
Isonitrosoguaiacol,  ethers  of,  in  their  relation  to  the  space  isomerism  of 

nitrogen  (Bridge  and  Morgan) II,  304 

Isomerism  (space)  of  nitrogen,  in  ethers  of  isonitrosoguaiacol  (Bridge 

and  Morgan)       II,  304 

in  ethers  of  toluquinoneoxime  (Bridge  and  Morgan)       11,145 

Isomerism  (space)  of  the  toluquinoneoxime  ethers  (Morgan)      ....  II,  283 

Laboratory  apparatus  (Gooch) I,  141 

Laboratory  apparatus  (Kreider) 1,306 


INDEX   OF  SUBJECTS.  411 

VOL.  PAGE 

Magnesium,  determination  of,  by  precipitation  as  ammonium  magne- 
sium phosphate  (Gooch  and  Austin)      11,190 

determination  of  barium  in  presence  of  (Mar) I,  125 

Maleic  anhydride,  action  of  primary  amines  upon  (Dnnlap  and  Phelps)  II,    44 

action  of  urea  and  primary  amines  upon  (Dunlap  and  Phelps)     .     .  II,    42 
Maleiiric  acid,  preparation  of,  from  urea  and  maleic  anhydride  (Dunlap 

andPhelps)          II,    42 

Manganese,  determination  of,  as  the  pyrophosphate  (Gooch  and  Austin)  II,  121 
condition  of    oxidation  of,  precipitated  by  the  chlorate  process 

(Gooch  and  Austin) II,    85 

estimation  of,  as  the  sulphate  and  as  the  oxides  (Gooch  and  Austin)  II,    77 

estimation  of,  separated  as  the  carbonate  (Austin) II,    96 

Manganous  chloride,  use  of,  in  hydrochloric  acid,  in  detection  of  oxidiz- 
ing agents  (Gooch  and  Gruener) I,  134 

use  of,  in  estimating  nitrates  (Gooch  and  Gruener) I,  132 

Mercury,  gravimetric  estimation  of,  as  the  oxalate  (Peters) II,  325 

separation  of  aluminum  from  ( Havens) 11,109 

titration  of,  by  sodium  thiosulphate  (Norton) II,  328 

Mercurous  oxalate,  separation  and  determination  of  (Peters)     ....  II,  320 
Metallic  salts,  action  of  sodium  thiosulphate  upon,  in  solution  at  high 

temperatures  and  pressures  (Norton) II,  384 

Molybdenum,  estimation  of  iodometrically  (Gooch) II,    27 

iodometric  determination  of  (Gooch  and  Norton) 11,111 

Molybdic  acid,  iodometric  estimation  of  (Gooch  and  Fairbanks)    ...  I,  375 

Monobromisonitrosoguaiacol  benzoyl  ether  (Morgan) II,  308 

Monobromtoluquinouemetaoxime  benzoyl  ether  ( Bridge  and  Morgan)    .  II,  1 58 

Monobromtoluquinoneorthooxime  benzoyl  ether  (Morgan) II,  286 

Naphthylmaleamic  acid  j8,  preparation  of,  from  0-naphthylamine  and 

maleic  anhydride  (Dunlap  and  Phelps) II,    45 

Nickel,  separation  of,  from  cobalt  (Havens) II,  141 

separation  from  cobalt  by  action  of  ammonium  hydroxide  on  the 

ferricyanides  (Browning  and  Hart  well) II,  344 

Nitrates,  action  of  phosphoric  acid  and  potassium  iodide  upon  (Gruener)  I,  193 

decomposition  of,  by  antimonious  chloride  (Gruener) I,  199 

detection  of  perchlorates  associated  with  (Gooch  and  Kreider)    .    .  I,  246 

iodometric  determination  of  (Gooch  and  Gruener) I,  132 

odometric  determination  of  (Gruener) I,  193 

Nitrates  and  chlorates,  estimation  of,  in  one  operation  (Roberts)  ...  I,  219 

Nitrates  and  nitrites,  estimation  of,  in  one  operation  (Roberts)      ...  I,  222 
Nitric  acid,  influence  of,  in  precipitation  of  barium  as  the  sulphate 

(Browning) I,  181 

reduction  of,  by  ferrous  salts  (Roberts) I,  203 

Nitric  oxide,  absorption  of,  by  iodic  acid  (Roberts)        I,  250 

Nitrogen,  space  isomerism  of,  and  bearing  of  ethers  of  toluquinoneoxime 

(Bridge  and  Morgan),  (Morgan) II,  145,  283 

space  isomerism  of,  in  ethers  of  isonitrosoguaiacol   (Bridge  and 

Morgan),  (Morgan) II,  304 

Nitrous  acid,  use  of,  in  liberating  iodine  (Gooch  and  Mar) I,    27 

use  of,  in  liberating  iodine  (Gooch  and  Ensign) I,    43 

Oxalic  acid,  titration  of,  by  potassium  permanganate  in  presence  of  hy- 
drochloric acid  (Gooch  and  Peters) II,  222 

Oxygen,  amount  of,  required  to  oxidize  an  organic  substance  (Phelps)    .  II,    71 


412  INDEX  OF  SUBJECTS. 

VOL.  PAGE 

determination,  In  air  and  in  aqueous  solution  (Kreider) II,    11 

Organic  acids,  application  of,  in  estimation  of  vanadium  (Browning  and 

Goodman)       II,      4 

Organic  substance,  amount  of  oxygen  required  for  oxidation  of  (Phelps)  II,    71 

combustion  of,  by  chromic  acid  in  the  wet  way  (Phelps)     ....  II,    67 

combustion  of,  by  potassium  permanganate  in  the  wet  way  (Phelps)  II,    62 

Perchlorates,  quantitative  determination  of  (Kreider) 1,316 

Perchloric    acid,  application  of,  to    the   determination  of    potassium 

(Kreider) I,  286 

preparation  of  (Kreider) I,  282 

Permanganate  solutions,  standardization  of,  by  electrolytic  iron  (Roberts)  I,  269 

Perofskite  (so-called),  analysis  of,  from  Magnet  Cove,  Ark.  (Mar)      .     .  I,    60 
Phosphorus,  iodometric    method    for    the   determination    of,  in    iron 

(Fairbanks) 1,391 

Phosphoric  acid,  determination  of,  by  precipitation  as  ammonium  mag- 
nesium phosphate  (Gooch  and  Austin)       II,  204 

Phosphoric  acid,  use  of,  with  potassium  iodide,  in  determining  nitrates 

(Gruener)        I,  193 

Phthalanil,  preparation  of,  from  phthalic  anhydride  and  sulphocar- 

banilide  (Dunlap) I,  361 

Phthalauilic  acid,  preparation  of,  from  phthalic  anhydride  and  sulpho- 

carbanilide  (Duiifap) 1, 361 

Phthalimide,  preparation  of,  by  action  of  urea  on  phthalic  anhydride 

(Dunlap) 1,355 

Primary  Amines,  action  of,  on  maleic  anhydride  (Dunlap  and  Phelps)    .  II,    44 
Potassium,  detection  and  determination  of,  spectroscopically  ( Gooch  and 

Hart) I,    92 

determination  of,  by  perchloric  acid  (Kreider) I,  282 

estimation  of,  as  the  pyrosulphate  (Browning) II,  368 

separation  of,  from  sodium  (Kreider  and  Breckenridge)      ....  1,401 
Potassium  Spectrum,  brightening  of,  by  sodium  chloride  (Gooch  and 

Hart) I,  101 

Potassium   bromide,  use  of,  in  reduction  of  selenic  acid  (Gooch  and 

Scoville) 1,335 

use  of,  in  reduction  of  arsenic  acid  (Gooch  and  Phelps)      ....  I,  265 
use  of,  with  hydrochloric  acid,  in  separation  of  arsenic  from  copper 

(Gooch  and  Moseley) I,  272 

Potassium  perchlorate,  decomposition  of,  by  anhydrous  zinc  chloride 

(Gooch  and  Kreider) I,  247 

Potassium  permanganate,  estimation  of  cerium  oxalate  by  (Browning 

and  Lynch) II,  297 

action  of  sulphuric  acid  upon  (Gooch  and  Danner) I,  145 

standardization  of,  in  iron  analysis  (Roberts) I,  269 

titration  of  oxalic  acid  by,  in  presence  of  hydrochloric  acid  (Gooch 

and  Peters) II,  222 

use  of,  in  combustion  of  organic  substances  in  the  wet  way  (Phelps)  II,    62 
use  of,  in  estimation  of  copper,  with  separation  from  cadmium, 

arsenic,  tin,  iron,  and  zinc  (Peters) II,  347 

use  of,  in  the  volumetric  estimation  of  mercury  as  the  oxalate 

(Peters) 11,320 

use  of,  in  the  estimation  of  selenious  acid  (Gooch  and  demons)  .    .  I,  297 

use  of,  in  the  estimation  of  tellurous  acid  (Gooch  and  Danner)    .    .  I,  154 


INDEX  OF  SUBJECTS.  413 

VOL.  PAGE 

use  of,  in  the  estimation  of  tellurous  acid  (Gooch  and  Peters)      .    .  II,  238 

Reducing  agents,  action  of,  on  iodic  acid  (Roberts) I,  250 

Rhodochrosite,  analysis  of,  from  Franklin  Furnace,  N.  J.  (Browning)    .  I,    57 

Rubidium,  estimation  of,  as  the  acid  sulphate  (Browning) II,  370 

quantitative  spectroscopic  determination  of  (Gooch  and  Phinney)    .  I,  157 

Selenic  acid,  iodometric  determination  of  (Gooch  and  Peirce)     ....  I,  341 

reduction  of,  by  hydrochloric  acid  (Gooch  and  Evans) I,  331 

reduction  of,  by  potassium  bromide  in  acid  solution  (Gooch  and 

Scoville) lt  335 

Selenious  acid,  determination  of,  by  potassium  permanganate  (Gooch 

and  demons) I?  297 

influence  of  hydrochloric  acid  in  thiosulphate  titrations  of  (Norton)  II,  206 

iodometric  determination  of  (Gooch  and  Peirce) 1,338 

Selenium,  gravimetric  determination  of  (Peirce) I,  365 

method  for  separation  of,  from  tellurium  (Gooch 'and  Peirce)  ...  I,  348 

reduction  of  acids  of,  by  hydriodic  acid  (Gooch  and  Reynolds)    .    .  I,  310 

Selenium  monoxide,  on  the  existence  of  (Peirce) 1,385 

Silver  salts,  electrolytic  reduction  of,  in  estimation  of  halogens  (Gooch 

and  Fairbanks)             I,  290 

Silver  sulphocyanide  in  gravimetric  analysis  (Van  Name) II,  359 

Sodium,  estimation  of,  as  the  pyrosulphate  (Browning) 11,371 

separation  of,  from  potassium  (Kreider  and  Breckenridge)      ...  I,  401 

Sodium  chloride,  brightening  of  potassium  spectrum  by  (Gooch  and  Hart)  I,  101 
Sodium  thiosulphate,  action  of,  on  solutions  of  metallic  salts  at  high 

temperatures  and  pressures  (Norton) II,  384 

influence  of  hydrochloric  acid  in  titrations  by,  with  special  reference 

to  the  estimation  of  selenious  acid  (Norton) II,  206 

reduction  of  iron  in  ferric  state  by  (Norton) II,  230 

use  of,  in  titration  of  mercury  (Norton) II,  328 

titration  of,  by  iodic  acid  (Walker) II,    52 

Sodium  tungstate,  use  of,  as  a  retainer  for  boric  acid  (Gooch  and  Jones)  II,  178 
Space  isomerism  of  nitrogen,  bearing  of  ethers  of  toluquiuoneoxime  on 

(Bridge  and  Morgan)  (Morgan) 11,145,283 

Spectroscopic  determination  of  potassium   (quantitative)    (Gooch  and 

Hart) I,  92 

Spectroscopic  determination    of    rubidium    (quantitative)  (Gooch  and 

Phinney) I,  157 

Standard  solutions  of  tartar  emetic,  stability  of  (Gruener) I,  216 

Standardization  of  potassium  permanganate  in  iron  analysis  (Roberts)  .  I,  269 

Starch,  blue  iodide  of  (Roberts) I,  236 

Starch    blue,  conditions  governing    formation   and   decomposition    of 

(Roberts)    . 1, 236 

Strontium,  estimation  of,  as  oxalate  (Peters) 11,374 

quantitative  separation  of,  from  calcium,  by  action  of  amyl  alcohol 

on  the  nitrates  (Browning)        I,  121 

separation  of  barium  from,  by  action  of  amyl  alcohol  on  the  brom- 
ides (Browning)       •    •  1,168 

separation  of,  from  calcium,  by  action  of  amyl  alcohol  on  the  nitrate 

(Browning) I,  121 

Succinanil,  preparation  of,  from  succinic  anhydride  and  sulphocarban- 

ilide  (Dunlap) 1,363 

Succinic  anhydride,  action  of  sulphocarbanilide  upon  (Dunlap)      ...  I,  363 


414  INDEX   OF  SUBJECTS. 

VOl.  PAGB 

Succinimide,  preparation  of,  from  urea  and  succinic  anhydride  (Dunlap)  I,  359 
Sulphates,  detection  of,  in  presence  of  sulphides,  sulphites,  and  thiosul- 

phates  ( Browning  and  Howe) 11,134 

Sulphides,  detection  of,  in  presence  of  sulphates,  sulphites,  and  thiosul- 

phates  (Browning  and  Howe) II,  134 

Sulphites,  detection  of,  in  presence  of  sulphides,  sulphates,  and  thiosul- 

phates  (Browning  and  Howe) II,  134 

Sulphocarbanilide,  action  of,  on  certain  acid  anhydrides  (Dunlap)      .    .  I,  355 

action  of,  on  phthalic  anhydride  (Dunlap) I,  355 

action  of,  on  succinic  anhydride  (Dunlap) 1,359 

Sulphocyanides  of  copper  and  silver  in  gravimetric  analysis  (Van  Name)  II,  359 
Sulphuric  acid,  action  of  potassium  permanganate  upon  (Gooch  and  Dan- 

ner) I,  145 

Tartar  emetic,  stability  of  standard  solutions  of  (Gruener) I,  216 

Tellurium,  determination  of,  by  precipitation  on  the  iodide  (Gooch  and 

Morgan) n,      1 

method  for  the  separation  of  selenium  from  (Gooch  and  Peirce)      .  I,  348 
Telluric  acid,  iodometric  method  for  the  estimation  of  (Gooch  and  How- 
land)      1, 277 

Tellurous  acid,  determination  of,  by  potassium  permanganate  (Gooch 

and  Danner) I,  154 

determination  of,  in  presence  of  haloid  salts  (Gooch  and  Peters)  .     .  II,  238 

Thallium,  estimation  of,  as  acid  and  neutral  sulphates  (Browning)    .    .  II,  317 

estimation  of,  as  chromate  (Browning  and  Hutchins) II,  300 

Thiosulphates,  detection  of,  in  presence  of  sulphides,  sulphites,  and  sul- 
phates (Browning  and  Howe) II,  134 

Tin,  detection  of  arsenic  associated  with  (Gooch  and  Hodge)      ....  I,  231 

separation  of  copper  as  oxalate  from  (Peters) H,  347 

Titanium  salt,  action  of  sodium  thiosulphate  upon,  at  high  temperatures 

and  pressures  (Norton) II,  392 

Toluquinoneoxime,  ethers  of,  and  their  bearing  on  the  space  isomerism 

of  nitrogen  (Bridge  and  Morgan) II,  145 

Toluquinoneoxime  ethers,  space  isomerism  of  (Morgan) II,  283 

Toluquinonemetaoxime  acetyl  ether  (Bridge  and  Morgan) 11,153 

Toluquinonemetaoxime  benzoyl  ether  (Bridge  and  Morgan) II,  154 

Toluquinonemetaoxime  methyl  ether  (Bridge  and  Morgan) II,  152 

Toluquinonemetaoxime,  sodium  salt  of  (Morgan) II,  285 

Toluquinoneorthooxime  acetyl  ether  (Bridge  and  Morgan) II,  160 

Toluquinoneorthooxime  benzoyl  ether  (Bridge  and  Morgan)      .    .    .    .  II,  160 

Toluquinoneorthooxime  benzoyl  ether  dichloride  (Morgan) II,  287 

Toluquinoneorthooxime  methyl  ether  (Bridge  and  Morgan)       .     .     .     .  II,  159 
Tolylmaleamic  acid  (o),  preparation    of,  from  maleic  anhydride  and 

o-toluidine  (Dunlap  and  Phelps)        II,  145 

Tolylmaleamic  acid   (p),  preparation  of,  from  maleic  anhydride  and 

p-toluidine  (Dunlap  and  Phelps) II,    44 

Urea,  action  of,  on  certain  acid  anhydrides  (Dunlap) I,  355 

action  of,  on  maleic  anhydride  (Dunlap  and  Phelps) II,    42 

Vanadic  acid,  reduction  of  by  hydriodic  and  hydrobromic  acids  (Brown- 
ing)      1,397 

estimation  of,  iodometrically  (Browning) 1,397 

Vanadium,  application  of  certain  organic  acids  to  estimation  of  (Brown- 
ing and  Goodman)       II,     4 


INDEX  OF  SUBJECTS.  415 

VOL.  PAGE 

Volatilization  of  the  iron  chlorides  in  analysis  (Gooch  and  Havens)    .    .  11,215 

Volumetric  estimation  of  mercury  (Peters) 11,320 

Zinc,  separation  of  aluminum  from  (Havens) II,  107 

separation  of  copper  as  oxalate  from  (Peters) II,  357 

Zinc  ammonium  phosphate  in  analysis  (Austin) II,  257 

Zinc  chloride  (anhydrous),  use  of,  in  detecting  perchlorates  (Gooch  and 

Kreider) 1, 247 

Zirconium,  separation  of   iron    from,  by  gaseous    hydrochloric  acid 

(Havens  and  Way) 11,266 

Zirconium  salt,  action  of  sodium  thiosulphate  upon,  at  high  temper- 
atures and  pressures  (Norton)       II,  391 


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THIS  BOOK  IS  DUE  ON  THE  LAST  DATE 
STAMPED  BELOW 


JAN  31  1916 


30m-l,'15 


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